KINETIC INVESTIGATIONS OF OXIDATION OF
S-PHENYLMERCAPTOACETIC ACIDS AND REACTIVITY OF
CYCLANOLS BY QUINOXALINIUM DICHROMATE
Thesis submitted to the
Bharathidasan University, Tiruchirappalli
for the award of the Degree of
DOCTOR OF PHILOSOPHY IN CHEMISTRY
by
G. MANIKANDAN, M.Sc., M.Phil.,
DEPARTMENT OF CHEMISTRY (DST-FIST Sponsored)
NATIONAL COLLEGE (AUTONOMOUS)
Nationally Reaccredited with „A‟ Grade by NAAC
College with Potential for Excellence by UGC
Tiruchirappalli 620 001
Tamilnadu, INDIA
MAY 2013
To
my Beloved Parents
Shri. A. Govindasamy, Smt. G. Indirani
Brothers and Sister
Dr. K. G. SEKAR, M.Sc., M.Phil., M.Ed., PGDCA, Ph.D.,
Associate Professor,
Department of Chemistry,
National College (Autonomous),
Tiruchirappalli - 620 001
Tamilnadu, India
Date:
-08-2014
CERTIFICATE
This
is
to
certify
that
the
thesis
entitled,
“KINETIC
INVESTIGATIONS OF OXIDATION OF S-PHENYLMERCAPTOACETIC
ACIDS AND REACTIVITY OF CYCLANOLS BY QUINOXALINIUM
DICHROMATE”, submitted to Bharathidasan University for the award of
the degree of Doctor of Philosophy in Chemistry is a bonafide record of the
research work done by Mr. G. MANIKANDAN under my guidance and
supervision during 2007-2013. This is also to certify that the thesis
represents the independent original work and has not previously formed the
basis for the award of any degree, diploma, associateship, fellowship or
other similar titles in any university.
(Dr. K. G. SEKAR)
Research Advisor
Mr. G. MANIKANDAN, M.Sc., M.Phil.,
Research Scholar (P.T.),
Department of Chemistry,
National College (Autonomous),
Tiruchirappalli - 620 001,
Tamilnadu, India.
Date: -08-2014
DECLARATION
The research work embodied in this thesis entitled “KINETIC
INVESTIGATIONS OF OXIDATION OF S-PHENYLMERCAPTOACETIC
ACIDS AND REACTIVITY OF CYCLANOLS BY QUINOXALINIUM
DICHROMATE”, is original and was done by me in the Department of
Chemistry, National College, Tiruchirappalli, Tamilnadu, India under the
guidance of Dr. K. G. SEKAR. It has not previously formed the basis for
the award of any degree, diploma, associateship, fellowship or other similar
title of Bharathidasan University or any other university.
(G. MANIKANDAN)
ACKNOWLEDGEMENT
First and foremost, I praise and thank from the depth of my heart, the
supreme power that has showered his benevolence in the completion of my
thesis work.
With immense pleasure, I express my deep sense of gratitude,
profound indebtedness and sincere thanks to my guide Dr. K. G. Sekar,
Associate
Professor,
Department
of
Chemistry,
National
College
(Autonomous), Tiruchirappalli for his inspiring guidance, invaluable
suggestions, constructive discussions and ceaseless encouragement during
every stage of research which was unique and I am fortunate in having
worked under his guidance. I am most grateful to him for the patience and
kindness shown to me during the tenure of my research.
It is my great pleasure to acknowledge with gratitude to
Shri. K. Raghunathan, Secretary, National College (Autonomous),
Tiruchirappalli for granting me permission to undertake fulltime research
work and providing the necessary facilities to complete the same.
I record my sincere thanks and gratitude to Dr. K. Anbarasu,
Principal, National College (Autonomous), Tiruchirappalli for allowing me
to carry out their research work peacefully in our campus.
I express my heartfelt thanks to Dr. K. Lakshmanan, Associate
Professor and Head, Department of Chemistry, National College
(Autonomous), Tiruchirappalli for providing the necessary facilities.
I feel proud in expressing my sincere thanks to the Doctoral
Committee member Dr. (Mrs.) P. Shameela Rajam, Associate Professor,
Department
of
Chemistry,
Bishop
Heber
College
(Autonomous),
Tiruchirappalli for her suggestions to complete my research work.
I have immense pleasure and take it my privilege to express my
heartfelt thanks to Dr. K.Anbarasu, Associate Professor of Chemistry, TRP
Engineering College (SRM), Tiruchirappalli, Dr. V.Palanivel, Associate
Professor of Chemistry, Periyar EVR College, Tiruchirappalli and
Dr. S.K.Periyasamy, Assistant Professor, Department of Chemistry, Jamal
Mohamed College (Autonomous), Tiruchirappalli.
I thank Dr. K. G.Sekar‟s family for their interest in my progress.
A million words would not be adequate for me to convey my gratitude
to my Friends.
On a personal ground, I am privileged to record a deep sense of
gratitude to my brothers Mr. G.Manoharan, Mr. G.Ravichandran and my
sister Ms. G.Kavitha for their support, love and encouragement for the
successful completion of the work.
(G. MANIKANDAN)
CONTENTS
Page
I.
INTRODUCTION
…
1
1.1
General Oxidation Reactions of Chromium (VI)
…
1
1.2
Oxidation States of Chromium
…
3
1.3
Kinetics of Chromic Acid Oxidation Reactions
…
4
1.4
Chromium (VI) as an oxidant
…
11
1.4.1 Chromic Acid
…
11
1.4.2 Pyridinium Bromochromate (PBC)
…
12
1.4.3 Pyridinium Chlorochromate (PCC)
…
12
1.4.4 2,2‟- Bi pyridinium Chlorochromate (BPCC)
…
15
1.4.5 Other Chromium (VI) Oxidants
…
16
1.5
Quinoxalinium Dichromate
…
17
1.6
S-Phenylmercaptoacetic Acids
…
19
1.7
Structure - Reactivity Relationships
…
25
1.8
The Hammett Equation
…
26
1.9
Cyclanols as Substrate
…
32
II.
SCOPE OF THE WORK
…
37
2.1
I-Strain Effect in Ring Compounds
…
38
III.
EXPERIMENTAL METHODS
…
39
3.1
Preparation of Quinoxalinium Dichromate
…
39
3.2
Preparation of S-Phenylmercaptoacetic Acid
…
39
…
40
3.2.1 Preparation of meta- methoxy
S-Phenylmercaptoacetic Acid
viii
3.2.2 Preparation of para- methoxy
…
40
…
41
…
41
…
42
…
42
…
43
…
43
…
44
…
44
…
46
acetic Acid by Quinoxalinium Dichromate
…
47
3.5
Non Kinetic study
…
49
3.6
Kinetic Measurements for Oxidation of Cyclanols
by Quinoxalinium Dichromate
…
51
Non Kinetic study
…
51
S-Phenylmercaptoacetic Acid
3.2.3 Preparation of meta-methyl
S-Phenylmercaptoacetic Acid
3.2.4 Preparation of para-methyl
S-Phenylmercaptoacetic Acid
3.2.5 Preparation of meta- bromo
S-Phenylmercaptoacetic Acid
3.2.6 Preparation of para- bromo
S-Phenylmercaptoacetic Acid
3.2.7 Preparation of meta- chloro
S-Phenylmercaptoacetic Acid
3.2.8 Preparation of para-chloro
S-Phenylmercaptoacetic Acid
3.2.9 Preparation of meta- nitro
S-Phenylmercaptoacetic Acid
3.2.10 Preparation of para- nitro
S-Phenylmercaptoacetic Acid
3.3
Determination of Physical Constants of Cyclanols
3.4
Kinetic Measurements for Oxidation of S-Phenylmercapto
3.7
ix
…
55
…
55
Acids by Quinoxalinium Dichromate
…
70
4.3
Correlation and Reactivity - Effect of Substituents
…
79
4.4
Kinetics and Mechanism of Oxidation of
in Aqueous Acetic acid Medium
…
81
4.5
Oxidation of Cyclanols by Quinoxalinium Dichromate
…
94
4.6
Effect of Structure on the Reactivity of Cyclanols
…
103
V.
SUMMARY
…
106
APPENDICES
…
109
A. Selected Values of Substituent Constant σ
…
109
B. Linear Regression Analysis
…
110
C. Symbols and Abbreviations
…
112
…
114
IV.
RESULTS AND DISCUSSION
4.1
Kinetics and Mechanism of Oxidation of
S-Phenylmercaptoacetic Acid by Quinoxalinium
Dichromate in Aqueous Acetic acid Medium
4.2
Oxidation of Substituted S-Phenylmercaptoacetic
Cylohexanol by Quinoxalinium Dichromate
REFERENCES
LIST OF PAPER PRESENTATIONS
LIST OF PAPER PUBLICATIONS
…
…
128
131
Chapter - I
INTRODUCTION
The kinetics of oxidation reactions and the investigation of the
reaction mechanisms from the kinetic data have been always the most
interesting subjects in chemistry. In any kinetic investigation, one may be
interested to arrive at (i) the relationship between the rate and the various
factors like concentrations of the reactants, temperature, reaction medium
etc., and (ii) interpretation of the empirical rate laws in the light of the
mechanism proposed.
The present study deals with these two aspects of certain selected
reactions, involving Chromium (VI) and its complexes as oxidants.
1.1
General Oxidation Reactions of Chromium (VI)1
Chromium trioxide is the most important chromium (VI) derivative. It
may be obtained on adding (i) sulphuric acid to an aqueous solution of
sodium (or) potassium dichromate (ii) on evaporating water from a reaction
mixture of potassium dichromate dehydrate and concentrated sulphuric acid.
The structure of chromium trioxide has been determined by X-ray analysis2
to be a linear polymer of chromium and oxygen atoms, with two additional
oxygen atoms linked to each chromium atom. Chromium trioxide dissolves
in water with accompanying polymerisation.
(CrO3)n + n H2O
n H2CrO4
The chromic acid formed is a fairly strong acid3.
… (1)
2
H2CrO4
H+ + HCrO4-
K1 = 1.21 mol/lit.
HCrO4-
H+ + CrO42-
K2 = 3.0 x 10-7mol/lit. … (3)
… (2)
In dilute aqueous solutions it largely exists as HCrO4-. In more
concentrated solution (> 0.05 M). It exclusively dehydrated to the
dichromate anion and its protonated forms4.
2HCrO4-
Cr2O72- + H2O
K3 = 35.5 mol/lit.
… (4)
HCr2O7-
H+ + Cr2O72-
K4 = 0.85 mol/lit.
… (5)
H2Cr2O7
H+ + HCr2O7-
K5 = large
… (6)
At still higher concentrations, poly chromates may be formed.
The foregoing equilibria are pH dependent. Above pH 8 only CrO 42ions exist. As the pH is lowered, i.e. between pH 2-6, HCrO4- and Cr2O72ions are at equilibrium. The equilibria are labile and by adding cations
insoluble chromates, only chromates are precipitated, no dichromates are
formed.
This kind of equilibria depends on the nature of the acid used. It
appears that there is a direct involvement of the mineral acid which
furnished the protons.
HCrO4- + 2H+
A-HCrO3A + H2O
… (7)
Thus the following species are formed during the protonation of
HCrO4- by the various mineral acids.
HCrO4- + H+ + H3PO4
HOCrO2 –OPO3H2 + H2O
… (8)
3
HCrO4- + H+ + HCl
HOCrO2 –Cl + H2O
… (9)
HCrO4- + H+ + H2SO4
HOCrO2– OSO3H + H2O
… (10)
It has also been suggested that Chromium (VI) in acetic acid may
exist in the form of an acetyl chromate ion5. Thus,
HCrO4- + HOCOCH3
O-–CrO3COCH3 + H2CrO4OCOCH3 … (11)
Complex formation between chromic acid and an anion results in a
charge in the dissociation constant. The more electron withdrawing power of
the anion of the mineral acid, the larger, the dissociation constant of the
complex species is given in terms of the mineral acids as6,
H3PO4 < HCl < H2SO4
Several salts of the type, M (CrO3X), have been described
(X = halogen, M = alkali metal or ammonium). For example, Potassium
chlorochromate can be prepared as orange crystals simply by dissolving
potassium dichromate in hot 6 M hydrochloric acid. By adding an organic
heterocyclic base to an aqueous solution of chloro chromic acid, the
corresponding salt may be obtained as coloured crystalline solid.
HCrO3Cl + B:
1.2
BH+CrO3Cl-
… (12)
Oxidation States of Chromium
The element exists in all oxidation states from 2 to 6 +, the highest state
(6+) corresponds to the sum of the 3d and 4s electrons analogous to titanium
and vanadium.
4
The most common and stable oxidation states are 2+, 3+ and 6+. The
2-, 1-, 0 and 1+ states are found in carbonyls, nitrosyls and in organometallic
complexes. The lowest oxidation states act as strong reducing agents. Thus
Cr2+, the first oxidation state known in aqueous solution, is widely used to
perform reductions both in organic and in inorganic reactions. The most
stable oxidation state is 3+. The oxidation states 4+ and 5+ are relatively rare.
Only a few compounds of Cr4+ and Cr5+ have been isolated and they
appear to be unstable in water as they rapidly disproportionate to Cr3+ and
Cr6+ compounds respectively. Cr4+ and Cr5+ species are however important
in Cr6+ induced oxidations.
1.3
Kinetics of Chromic Acid Oxidation Reactions
Although, a variety of compounds can be oxidised by chromic acid,
Dash et al.7 used chromic acid as an oximetric titrant. The oxidation of
malonic acid by acid dichromate in aqueous acetic acid medium has been
found to be first order with respect to both oxidant and substrate involving
complex formation between them. The acid dependence shows first order in
(sulphuric acid) and beyond two molar acid concentrations, the order was
found to be more than 3.0. The participation of both HCrO4- and Cr2O72- has
been inferred.
Kinetics of chromic acid oxidation of substituted mandelic acids was
followed by Sundaram and Venkata Subramaniyan8 in acetic acid - water.
The effect of various substituents and the structural influences has been
analysed, with the help of Hammett equation.
5
Venkataraman and Brahamaji Rao9 followed the oxidation kinetics of
formic acid by chromic acid, at different molar concentrations of sulphuric
acid, ranging from 1 to 5.5. The data have been examined in the light of
Bunnett‟s criteria of reaction mechanism. A tentative mechanism involving
both SN1 and SN2 reaction at lower pH was proposed.
The same kinetics when followed by Obula Reddy and Brahamaji
Rao10 in moderately high concentrations of phosphoric acid (1.0 to 7.0 M), a
pronounced rate enhancement was observed. There was a linear
proportionality between the rates and the concentrations of formic acid. This
observation can be employed for the analytical determination of even small
amount of formic acid.
Sen Gupta et al.11 made an extensive kinetic study on the oxidation of
α-hydroxy isobutyric dl-α-phenyl acetic acid and citric acids. The kinetic
results do not indicate the formation of the intermediate compound between
Cr (VI) and the substrates. The reactions also do not proceed via free
radicals. A mechanism based on the formation of carbonium ion in the slow
rate determining step has been suggested.
Singh Dhakaray and Ghosh12 followed the acid catalysed oxidation of
mandelic acid in the pH range, 1.90-3.30. Bivalent manganese accelerates
the rate of the reaction. The probable mechanisms for Mn2+ catalysed and
uncatalysed reactions were discussed.
Paul and Pradhan13 proposed a novel mechanism (Scheme 1.1) for the
oxidative decarboxylation of mandelic acid based on a study of deuterium
labelling, solvent isotope effect.
6
H
H
C6H5 C OH + HO
O
Cr
O C OH
C6H5 C O
Cr
-2H2O
HO
O
+
slow
O
O
C O
O
H
C6H5 C O
O C O CrO2-
+
C6H5 C OH
O C O CrO2-
+
C6H5 C O
H + CrIV
C6H5 C O + CO2
O C O
H
Scheme 1.1
The mechanism shown is proposed for decarboxylation wherein a
cyclic anhydride is formed first. It then breaks up followed by the hydride
ion transfer from -carbon to the adjacent electron-deficient oxygen. This is
likely as the positive charge on the -carbon atom can be stabilized by the
-electron cloud of the benzene ring.
Spectrophotometric method was caused to follow the kinetics of
oxidation of thioglycollic, thiolactic and thiomalic acids, in sodium acetate acetic acid buffer14. The formation of free radical intermediate has been
demonstrated during the oxidation of thioglycollic acid, whereas there was
no experimental evidence for the free radical mechanism, in the case of thio
lactic and thiomalic acids.
7
The kinetics of oxidation of glycollic acid by chromium peroxy
dichromate was followed by Valachha and Dakwale15 at low concentrations
of the substrate, the order of the reaction was pseudo-zero and one with
respect to oxidant. At high concentrations of the substrate, the respective
orders were found to be one and two. The product of oxidation was
formaldehyde. Influence of temperature, solvent, acids and added salts of
Mn (II) and Cr (III) has been studied.
Singh et al.16 has reported the result of the process of oxidation of
some hydroxy acids by Cr (VI).
Acid dichromate oxidation kinetics of a few α-hydroxy acids such as
mandelic, lactic, benzillic and benzyl phenylglycollic acids by Saran and
Acharya17 showed that the reaction followed first order each with respect to
oxidant and every substrate studied.
R1
R
-
C O
+
RR1C(OH)COOH + HCrO4 + H
O
Cr
O
C O
+ 2H2O
O
R1
R
C O
O
slow
Cr
O
C O
RR1C O + CO2 + Cr4+
O
Cr4+ + Cr6+
fast
2RR1C(OH)COOH + 2Cr5+
2Cr5+
fast
Scheme 1.2
2RR1C O + 2CO2 + 2Cr3+
8
The reaction was acid catalysed and the order was unity in (H+). Both
HCrO4- and Cr2O72- seemed to participate in the reaction consistent with the
result of oxidation and product analysis, a reaction path involving the
formation of cyclic chromate ester (Scheme 1.2), decomposing the rate
limiting step by oxidative decarboxylation has been suggested. The
oxidation was proceed by a two electron transfer.
Radhakrishna moorthy and Pande18 made an extensive kinetic study of
Os (VIII) catalysed chromic acid oxidation of maleic, fumaric, acrylic and
cinnamic acids in aqueous and in aqueous acetic acid media, in the presence
of perchloric acid. Maleic acid and cinnamic acids exhibit zero order
dependence in oxidant at lower concentration of oxidant and first order
dependence at higher concentration of oxidant. Acrylic acid showed zero
order dependence in oxidant in the total range of [oxidant] studied, while
fractional order dependence in oxidant was noted in the case of fumaric acid.
The order in substrate was unity in all the case of study. The effect of acidity
was marginal and the rate decreased slowly with the increase of percentage
of acetic acid.
Chromic acid oxidation of aromatic acetals (got from benzaldehyde
and aliphatic alcohols) studied by Nambi et al.19 in aqueous acetic acid,
yielded the corresponding esters as the main products. A total second order
kinetics, first order each in [acetal] and [Cr (VI)] was observed correlation
analysis of the rate data, elimination of proton in the rate determining step
were the prominent points of this study.
Kinetics of chromic acid oxidation of dimethyl malonate by Oswal20
in acetic acid - water solutions of H2SO4 - H3PO4, in the presence and in the
9
absence of Mn (II) ions, was studied at constant ionic strength. A welldefined induction period, marked catalysed activity of Mn (II) species and
the thermodynamic quantities of uncatalysed reaction were considered to
propose a probable mechanism involving free radicals.
A mechanism involving the formation of an iminoxy radical in the
rate determining step was proposed during the kinetics of oxidation of some
para- substituted acetophenone oximes by Cr (VI)21. Oxidative hydrolysis of
the reaction showed a first order rate dependence on the substrate
concentration but inverse dependence on the concentration of Cr (VI). The
ion-dipole type of this reaction was favoured by electron donating
substituents with a reaction constant δ = -0.7.
Several substituted N-methyl-2,6-diphenyl piperidin-4-ones22 were
subjected to oxidation by aqueous acidic CrO3, to investigate the effect of
3-alkyl substituent on this reaction. Increase of rate with increase of [H +],
solvent composition and ionic strength (due to the addition of Na2SO4) was
observed. A suitable mechanism involving a rate determining formation of a
chromate ester between CrO3 and piperidone was discussed in detail.
Kinetics of oxidation of aliphatic acetals23 (prepared from aliphatic
aldehydes, aliphatic alcohols, halogen substituted alcohols and aromatic
alcohols) by chromic acid in acetic acid medium showed first order each in
oxidant and acetal the corresponding ester was the main product. Substituent
effect, activation parameters and salt effect suggested that the elimination of
a proton from the complex species involving the acetal and chromium,
would be the rate determining step.
10
It was shown that the oxidation of dipentyl and diphenyl sulphoxides
followed with Cr (VI)24 in sulphuric acid medium involved an electron
transfer from the sulphoxide to Cr (VI) and HCrO3- in the rate determining
step. A cation radical rapidly attacking the Cr-O bond gave rise to a
complex, undergoing subsequent hydrolysis to yield the corresponding
sulphone, camphor when subjected to Cr (VI) oxidation25 showed a total
second order kinetics, the condition of constant acidity. Addition of Mn (II)
and Co (II) ions retarded the rate while ethylene diamine facilitated the
same. A suitable mechanism has been proposed.
Kinetics of oxidation of diethyl tartarate26 by chromic acid was found
to be first order each with respect to Cr (VI), ester and H+ ion. The product
of oxidation was ethyl glyoxalate. The results of the overall second order
kinetics of oxidation of some aliphatic aldehydes by chromic acid 27 were
discussed in the light of the theories of Amis and Laidler.
Electron releasing substituents enhanced the rate of oxidation of parasubstituted toluenes by CrO328 in acetic acid - water mixtures at [HCrO4-]
0.1 to 1.0 M. Radical intermediates were formulated to propose a suitable
mechanism, to compute the δ value and to explain the importance of acidity
function OH, rather than H+.
Ramanathan and Varadarajan29 studied the kinetics of oxidation of
benzoin by chromic acid. The rate was proportional to the first power of
concentration of each of benzoin and Cr (VI). The authors reported that the
rate determining enolization envisaged in other cases of similar study is not
consistent with their results.
11
1.4
Chromium (VI) as an Oxidant
Chromium compounds have been widely used in aqueous and non-
aqueous medium for1, 30-35 the oxidation of a variety of organic compounds,
chromium compounds especially Cr (VI) reagents have been proved to be
versatile reagents capable of oxidising almost all the oxidisable organic
functional groups1, 36-40.
Generally Cr (VI) oxidation reactions have been performed in
aqueous acidic conditions, the source of Cr (VI) being chromium trioxide
and sodium or potassium dichromate. Aqueous sulphuric acid and acetic
acid are the most frequently employed acids, acetone, benzene, methylene
chloride or ether are often used as co-solvents in the case of water insoluble
organic compounds.
However, a variety of new chromium (VI) oxidants together with
special reaction conditions have been introduced for the chemospecific,
regiospecific and stereospecific oxidative degeneration of functional groups
in highly sensitive systems.
1.4.1 Chromic Acid
Chromic acid exhibits in aqueous solution in the absence of other
ions41 viz., HCrO4-, CrO42- and Cr2O72-. The dichromate ion and its
protonated forms are the predominant species at concentrations higher than
0.05 M. Chromium (VI) is mainly in the monomeric form in very dilute
solutions. Only in the basic solution
(pH > 7) acid chromate ion, HCrO4-
loses a proton. It gains a second proton in the pH range of 1 to 3.
12
According to Wiberg and Schaffer42 chromium (VI) exists in 97%
acetic acid as the aceto chromate ion (CH3COCrO2O-) as the principal anion
in the oxidation of propan-2-ol by chromic acid. Rocek and Krupicka43 have
postulated HCrO3- and H3CrO4+ species in strongly acidic media.
In presence of added ions like Cl-, Br-, F-, SO42- and PO43- new
chromium species have been formed in aqueous medium4.
O
X- + H+ + HO
Cr
O
O-
X Cr
O
O- + H2O
… (13)
O
1.4.2 Pyridinium Bromochromate (PBC)
Narayanan and Balasubramaniyan44 have found PBC as an efficient
oxidant for alcohols as well as a brominating agent for aromatic compounds.
1.4.3 Pyridinium Chlorochromate (PCC)
Banerji et al.45 studied the kinetics of oxidation of thioglycollic acid,
thiolactic acid and thiomalic acid by PCC. The reaction is first order with
respect to [PCC] and Michaelis-Menten types of kinetics were observed
with respect to all the [thioacids]. The rate was not affected by the addition
of acrylonitrile indicates the absence of the free radical mechanism. From
the results, a suitable mechanism was proposed as follows.
13
O
O PyH+
R S H + Cr
O
O
K
Cr O PyH+
R S
Cl
HO
Cl
K2
R S+ + [HOCrOClOPyH]R S+ + R S H fast
R S S R + H+
[HOCrOClOPyH]- + H+ fast H2O + CrOClOPyH
Scheme 1.3
Kinetics of oxidation of some substituted phenyl methyl sulphides to
their
corresponding
sulfoxides
by
PCC
have
been
studied
by
Rajasekaran et al.46 in binary solvent moistures of 60% (v/v) aqueous acetic
acid and 50% (v/v) chlorobenzene nitrobenzene. The reaction in dipolar
protic solvent systems is strictly second order, while in aprotic solvents it
follows Michaelis-Menten kinetics. Hammett ρ value (-2.12 ± 0.09)
indicates an electron deficient transition state.
14
C6H5
O-Py+H
O
: S: + Cr
CH3
O
C6H5
S….+….Cr
K
Cl
O-Py+H
O
CH3
O
Cl
(C1)
C6H5
O-Py+H
O
K1 C6H5SCH3 + CrO2 + Py+HCl-
S………Cr
CH3
O
O
Cl
(C1)
Scheme 1.4
The oxidation of phosphinic, phenyl phosphinic, phosphonic acids
with PCC have been studied by Banerji et al.47. Ketones have been
synthesised from organo boranes using PCC by Parish et al.48. PCC oxidises
D-ribose to D-riboxyacetone49. Jha and Agarwal50 carried out the kinetics of
oxidation of pinacol in the nitrobenzene - methylene chloride medium.
The oxidation of some para- and meta- substituted benzaldehydes by
PCC51 is first order each in [substrate], [PCC] and [H+]. The order of the
reactivity is given as p-NO2 > p-COOEt > p-I > p-Br > p-Cl > -H > p-CH3 >
p-OCH3 > m-NO2 > m-Br > m-Cl > m-I > m-OCH3 > m-CH3. Electronreleasing substituents retard and the electron-withdrawing groups enhance
15
the rate. The change in ionic strength has no effect. An initial two electron
transfer has been formulated and chromium (VI) thus formed reacts with
aldehyde to give chromium (III) and aryl radicals which is tested by
polymerisation of acrylonitrile and EPR spectrum. The reaction proceeds
with the formation of a PCC ester as the intermediate.
The oxidation kinetics of glucose, with PCC in perchloric acid
medium has been studied by Dhar52. Ananthakrishna Nadar et al.53 studied
the kinetics of oxidation of substituted styryl phenyl ketones and of
substituted styryl methyl ketones by PCC in 90% acetic acid and 10%
water (v/v) containing perchloric acid and NaClO4 at 30 0C, 50 0C and 60 0C.
The two reactions are the first order each in ketone and PCC. The rate
constants are well correlated with σ+ constants. The effects of varying
[HClO4] and the percentage of acetic acid on the reaction rate have also been
studied.
1.4.4 2,2‟ - Bipyridinium Chlorochromate (BPCC)
BPCC is a useful oxidising agent for the conversion of primary and
secondary alcohols to carbonyl compounds. Its use simplifies the
purification of the resulting carbonyl compound.
This reagent because of the following characteristic properties can be
used as a good oxidising agent both in kinetic as well as in synthetic
reactions54.
a. It is soluble in non-aqueous solvents and aqueous solvents.
16
b. It is yellow crystalline non hygroscopic and a stable and still
effective after three months of storage.
c. It liberates iodine instantaneously from potassium iodide solution.
Kabilan et al.55 have determined the effect of ring size on the rate of
oxidation of cyclanols by BPCC in acetonitrile medium.
1.4.5 Other Chromium (VI) Oxidants
Chromium trioxides associated with some basic ligands leading to
various new types of Cr (VI) oxidants are also available. Such compounds
are tetra butyl ammonium chromate (Bu4N)2CrO4, tetra butyl ammonium
dichromate
(Bu4N)2Cr2O7,
tetra
butyl
ammonium
tetra
chromate
(Bu4N)2Cr4O13 etc.,
Some of the other Cr (VI) oxidants reported in the literature are as follows:
2,2‟ – Bipyridininium Chromate (BPC)56
4 – (dimethyl amino) Pyridinium Chlorochromate57
Imidazolium Dichromate58
Nicotinium Dichromate59
Phthalazinium Chlorochromate (PtCC)60
Phthalazinium Dichromate (PtDC) 60
Pyrazinium Dichromate (PzDC) 60
Pyrazinium Chlorochromate (PzCC)61
Quinolinium Chlorochromate62
Quinolinium Dichromate62
17
1.5
Quinoxalinium Dichromate [(C8H6N2+H2)Cr2O72-]
The chromium based reagents developed so far all suffer from at least
one of the following drawbacks.
i. Highly acidity.
ii. Photosensitivity.
iii. Instability.
iv. Tedious workup procedures or requirement of large excess of
reagent.
Therefore, the search for a new reagent persisted which has now led to
the synthesis of quinoxalinium dichromate. The possible structure of QxDC
is,
H
N
Cr2O72N
H
Quinoxalinium dichromate (QxDC)
Quinoxalinium dichromate (QxDC)63 can be easily prepared in good
yield (78%) by addition of quinoxaline to a solution of chromium trioxide in
water in a molar ratio of 1:1. QxDC is a yellow, non-hygroscopic and stable
solid compound which can be stored in the darkness for months without
losing its activity. The structure of the product was confirmed by elemental
analysis and its IR spectrum. In order to ascertain the efficiency of the
18
reagent as an oxidant, it was tested on a wide array of substrates in
dichloromethane at room temperature. Moreover, it is stable and can be
stored for long periods without much loss in its activity and hence turns out
to be a very useful reagent in synthetic organic chemistry.
Oxidation of some primary and secondary alcohols by quinoxalinium
dichromate was studied by Degirmenbasi63. In this study, oxidants were
carried out in dichloromethane with a substrate to oxidant ratio of 1:1.5
at room temperature. The products of the reactions were corresponding
aldehydes and ketones, identified by comparison of their physical and
spectroscopic data with those of authentic samples in the presence of
anhydrous acetic acid as catalyst.
Ozgun64 studied the oxidation of substituted benzyl alcohols by
quinoxalinium dichromate. A kinetic study quinoxalinium dichromate
oxidizes benzyl alcohol and substituted benzyl alcohols smoothly in
dimethyl sulfoxide and in the presence of acid to the corresponding
aldehydes. The reaction has unit dependence on each of the alcohol, QxDC
and acid concentration. Electron-releasing substituents accelerate the
reaction, whereas electron-withdrawing groups retard the reaction and the
rate data obey Hammett‟s relationship. The analysis of the dependence of
the kinetic isotope effect on temperature indicated that the reaction involves
a symmetrical cyclic transition state. The rates of oxidation were determined
at different temperature and the activation parameters were evaluated.
A suitable mechanism is proposed.
19
1.6
S-Phenylmercaptoacetic Acids
Oxidation of S-phenylmercaptoacetic acid is interesting in the fact
that it can undergo a Pummerer type of rearrangement followed by the
cleavage of the molecule leading to the products thiophenol and glyoxalic
acid65-74. The rearrangement takes place due to the instability of the
intermediate, α-sulfinyl acetic acid in acetic medium. Similarly α-sulfinyl
ketones and β-disulfoxides are also unstable in acidic conditions.
The instability of sulfoxide in the presence of acid varying from dilute
mineral acids through dry hydrogen halides to mercuric chloride has been
reported earlier75-79.
Generally oxidation of organic sulphides by various oxidising
reagents leads to either sulfoxide or sulfone depending on the reaction
conditions. However the oxidation of S-phenylmercaptoacetic acid differs
from that of alkyl or aryl sulphides due to the presence of an active
methylene group adjacent to the sulfur atom. Though, the product of
oxidation is phenyl sulfinyl acetic acid, the instability of the same leads to
the rearrangement in presence of acids.
The mechanism given by pummerer65,
66
for the acid catalysed
oxidative cleavage of phenyl sulfinyl acetic acid is given below
(Scheme 1.5)
20
O
OH...H
C6H5SCH2CO2H
H+ C6H5 S
CHCO2H
(I)
OH
+
OH2
+H
C6H5SCHCO2H
+
-H
C6H5SCHCO2H
(II)
C6H5SH + HCOCO2H
Scheme 1.5
Later Kenney, Walsh and Davenport put forward the following
mechanism (Scheme 1.6):
O
C6H5SCH2CO2H
O
H+ C6H5
+
S CH2CO2H
H
OH
C6H5SCHCO2H -H+ C6H5 S O CHCO2H
H
C6H5SH + HCOCO2H
Scheme 1.6
H
21
The mechanism suggested by Pummerer has been supported by
further evidence given by Walker and Leib68. Also a new interpretation of
the mechanism of the acid-catalysed cleavage of phenyl sulfinyl acetic acid
has been proposed. The scheme provided by Walker and Leib is as follows:
O
OH...H
C6H5SCH2CO2H H+ C6H5 S
CHCO2H
(I)
OH
OH2
C6H5SCHCO2H
+H+ C6H5SCHCO2H
-H+
(II)
C6H5SH + OHCCO2H
Scheme 1.7
It has been suggested that the above reaction may be concerted in that
loss of proton from the methylene carbon atom could coincide with the
rearrangement as indicated in I.
The rapid oxidative cleavage of the carboxy-methyl group suggested
the utility of this group as a readily removable sulfur protective species
which would enable electrophilic substitution of the aromatic ring of
thiophenols. Subsequently a number of substituted thiophenols have been
prepared using this reaction67, 69. All these reactions were shown to proceed
through the formation of phenyl sulfinylacetic acid intermediate.
22
The oxidants used to cleave the molecule are mainly hydrogen peroxide,
nitric acid and permanganate.
Kenney,
Walsh
and
Devenport
have
made
the
following
generalizations regarding this reaction.
i.
α-sulfinyl acids, α-sulfinyl esters, α-sulfinyl ketones and
β-disulfoxides disproportionate under a wide variety of acidic
conditions to give products in which the sulfur atom has been
reduced and the α-carbon atom oxidized.
ii.
Acid catalysis is a necessary factor.
iii.
For the disproportionation to take place, the carbon atom α-to
the sulfoxide must bear a hydrogen atom.
iv.
When the α-carbon bears a strong electron withdrawing group,
the reactions is greatly facilitated.
v.
The presence of a substituent like p-CH3 group in benzene ring
of phenyl sulfinyl acetic acid promotes the disproportionation,
whereas a p-NO2 group retards it.
Though this reaction has been well established by several possible
mechanisms, it seems that it has not yet been investigated in detail through
kinetic studies. However, few reports are available on the kinetics of
oxidation of S-phenylmercaptoacetic acids.
23
Initially Srinivasan and Pitchumani have studied the kinetics of
oxidation of S-phenylmercaptoacetic acid using the oxidants chloramine-T80
and potassium peroxy disulphate81.
Kabilan et al.82 studied the oxidation of S-phenylmercaptoacetic acid
and phenoxy acetic acid by pyridinium dichromate. The reaction for
phenylmercaptoacetic acid is conducted in presence of oxalic acid, it acts as
a catalyst and also a co-substrate. The reaction for phenylmercaptoacetic
acid is conducted in presence of perchloric acid. Both the reactions have
been found to be acid catalysed one. The order with respect to PDC is one.
The reaction follows a Michaelis-Menten type of kinetics with respect to
substrate.
A plausible mechanism which is applicable to both the oxidation
reaction has been proposed. In aqueous acetic acid medium the effective
oxidizing species of a chromium (VI) reagent is reported to the HCrO4- ion.
Initially, the HCrO4- ions form a complex with the substrate in an
equilibrium step which is followed by the dissociation of the complex in
presence of H+ ions in a slow and rate determining step.
Oxidation cleavage of S-phenylmercaptoacetic acids by pyridinium
chlorochromate – kinetic and correlation analysis done by Kabilan et al83.
Oxidation of 24 S-arylmercapto acetic acid by pyridinium chlorochromate
have been studied in acid medium. The rate data of meta- and parasubstituted acids have been correlated well with σI, σR0 values and the metacompounds correlate well with F, R values.
24
Further, the ortho- substituted acids show a good correlation with
triparametric equation involving Taft‟s σI and σR0 and Charton‟s steric
parameter γ. There is no considerable steric contribution to the total ortho
substituent effect.
Sathiyanarayanan et al.84 studied oxidation of S-phenylmercaptoacetic
acid by N-Chloronicotinamide. A kinetic and mechanistic study, the
conversion of S-phenylmercaptoacetic acid to the corresponding sulfoxide
was performed in 50% (v/v) water-acetic acid mixture in the presence of
perchloric acid. The orders with respect to substrate, oxidant and perchloric
acid were one. Ionic strength had a considerable influence on the rate of the
reaction. It indicates the involvement of a dipole in the rate limiting step.
A suitable mechanism is consonance with the observed facts is proposed.
Kinetics
and
mechanism of
EDTA
catalysed
oxidation
of
S-phenylmercaptoacetic acid by chromium (VI) has been studied by
Sathiyanarayanan et al.85. The conversion of S-phenylmercaptoacetic acid to
the corresponding sulfoxide was performed in 50% (v/v) water-acetic acid
mixture in the presence of the disodium salt of ethylene diamine tetra acetic
acid, the catalyst, the ionic strength had no appreciable effect on the reaction
rate. The ratio K(D2O) / K(H2O) < 1 clearly indicates a significant solvent
isotope effect. Highly negative entropy (ΔS#) values indicate a structured
transition state. A mechanism is proposed involving the formation of a
ternary
complex
comprising
EDTA,
chromium
(VI)
S-phenylmercaptoacetic acid in a fast step. The complex hydrolyzes in a
and
25
subsequent slow rate-determining step yielding the sulfoxide. Electronreleasing substituents in the phenyl ring accelerate the rate, while electronwithdrawing substituents retard the rate.
Alhaji et al.86 studied mechanism of oxidation para- substituted
phenylthioacetic acids with N-Bromophthalimide (NBP). The kinetics of
oxidation of phenylthioacetic acid by N-Bromophthalimide in acetonitrilewater solvent mixture at 298 K in the presence of perchloric acid has been
followed potentiometrically. The reaction is first order with respect to
oxidant and substrate and inverse fraction order in H+. This reaction is in
favour of a SN2 type mechanism, involving NBP itself as the reactive
species. The electron-releasing substituent in the phenyl ring of
phenylthioacetic acid accelerates the reaction rate while the electronwithdrawing substituent retards the rate. The excellently linear Hammett plot
yields a larger negative ρ value supporting the involvement a bromo
sulphonium ion intermediate in the rate determining step.
1.7
Structure - Reactivity Relationships
The structural factors controlling the reactions may be obtained by the
application of the linear free-energy relationships. A substituent can
influence a distant reaction centre atleast by five different processes87.
i.
The electric dipole field of the polar substituent - substrate bond
can be influence the reaction centre across space.
ii.
The primary inductive effect can be transmitted to the reaction
centre by successive polarisation of the intervening sigma
bonds.
26
iii.
The electrostatic charge set up at a conjugate atom adjacent to
the substituent may polarise the corresponding п-electron
system.
iv.
The п-electron system can be polarised by resonance interaction
with the substituent.
v.
The electrometric effect, which has importance when there is
mutual conjugation between the substituent and the reaction
centre.
The importance of these various effects can be understood from many
examples87-91.
1.8
The Hammett Equation
Various linear free-energy relationships were discovered in the early
1930s for the side-chain reaction of meta- or para- substituted benzene
derivatives. Hammett‟s92 contribution in 1937 lay essentially in recognizing
the value of taking one reaction as a standard process, with which all other
relevant reactions could be compared. In terms of a very simple
mathematical equation much information about reactivity could be
summarised. Equations (14) and (15) show the basic forms of the Hammett
equation, in which K or k is the rate or equilibrium for a side chain reaction
of meta- or para- substituted benzene derivative (Fig. 1.1).
log K = log K0 + ρσ
… (14)
log k = log k0 + ρσ
… (15)
27
R - Substituent
Y - Reaction Centre
R
Y
(Fig. 1.1)
The symbol K0 (or) k0 value of rate or equilibrium constant for the
unsubstituted compound the substituent constant σ measures the polar
(electronic) effect of replacing H by a given substituent (in the meta- or
para- position) and is in principle, independent of the nature of the reaction.
The reaction constant ρ depends on the nature of the reaction and measure
the susceptibility of the reaction to polar effects. Hammett chooses
ionisation of benzoic acids in water at 25 0C as the standard process. For,
this ρ was taken as 1.0 arbitrarily and the value of ρ for a given substituent
then becomes log (ka/k0a), where ka is the ionisation constant of the
substituted benzoic acid and k0a that of benzoic acid itself.
When log K or log k as appropriate, is plotted against σ of meta- or
para- substituted compounds (for parent σ = 0) a straight line should be
obtained, however, by the method of least squares log K or log k is taken as
the explanatory variable. Jaffe93 examined its application to about 400
reaction series, losing great stress on the correlation coefficient (r) and the
Standard deviation (Sd) as a measure of success of the Hammett equation.
The parent of all such relationships of this kind was the discovery by
Bronsted and Pederson94 of the general acid base catalysis and at the
catalysed reactions that are linearly related to those of the acidity constants
28
of the catalysing acid or base. Pederson95 clearly recognised that this is a
relationship between the rate and the equilibria of the same series of
reactions, (i.e.,) proton transfer process. Hammett and Pfluger96 extended the
idea of finding out a quantitative relation between the logarithms of the rate
constant of reactions.
Burkhardt97 and Hammett98 found mainly Linear Free Energy
Relaionships (LFER) in the reactions of substituted benzene derivatives.
The equations (14 and 15) is often called as LFER and implies that
there is a linear relationship between free energies of activation for one
homologous series of reactions and those for another. This can be shown by
transforming the Hammett equation into the form with the help of Eyring‟s
equation99.
- ΔG#/RT = -ΔG0#/RT + ρσ
… (16)
In equivalent terms, equation (16) can be represented as follows.
- ΔG# = ΔG0# + ρσ RT
… (17)
Here ΔG0#, ρ and T are constants.
For a second reaction series having the reaction constant ρ‟.
-ΔG‟# = ΔG‟0# + ρσ RT
… (18)
Equation (17) and (18) then lead to the relationship.
ΔG#/ρ - ΔG‟#/ρ‟=ΔG0#/ρ - ΔG‟0#/ρ‟
… (19)
29
which may be written as
ΔG#= ρ/ρ‟ΔG‟# + constant
… (20)
Equation (20)100 implies the linear free energy relationship embedded
in the Hammett Equation, because it is in the form a linear equation and by
which the free energy changes associated with the members of one reaction
series can linearly related to the free energy changes associated with the
corresponding substituted members of another reaction series many
books101-107 and reviews108-115 are available in the subject of LFER.
The reaction constant ρ is a measure of the susceptibility of the
reaction centre to the influence of the substituents. A reaction which is
facilitated by reducing the electron density at the reaction centre has a
negative value.
Electron withdrawing substituents have positive values of σ and
electron releasing substituents have negative values. Hammett σ values of
the some common substituents are taken from Mc Daniel and Brown 116. The
sole effects of a substituent that alter the reactivities have been taken as
steric, inductive and resonance effects117. In the case of meta- and parasubstituted benzene derivatives there are no steric effects118.
The inductive effect is proportional to the inductive substituent
constant (σI)119. These σI values are based upon inductive effects from the
aliphatic series and the ionisation constant values of meta- and parasubstituted benzoic acids from the knowledge of σI, a set of σR values have
been derived. Therefore, σ = σI + σR where σR is the substituent contribution
30
through resonance with the benzene п-bonds to the electron density of the
nuclear carbon atom at which the functional group is attached. The
resonance contribution of the meta- substituent is approximately one third of
the corresponding para- substituent. This is due to the indirect resonance
interaction120, of the meta- substituent with the positions ortho- to the atom
bearing the side chain reaction centre as represented in (Fig.1.2 and Fig.1.3)
±
Y
Y
±
±R
(Fig. 1.2)
R
(Fig. 1.3)
Since the resonance of a functional centre with the benzene ring does
not produce resonance effects on rates in the absence of direct conjugation,
the effects of substituents are inductive-polar and resonance-polar. One may
therefore, take (σp-σm) as a measure of the resonance-polar effect of a
substituent. This difference generally gives the light quantitative order of
resonance-polar effects but does not give the right magnitude because the
inductive effect at meta- and para- positions are not identical121 and there is
relayed resonance effect of the meta- substituent from the ortho- positions to
the reaction centre. Thus it is clear that only (σ-σ‟) measures the resonancepolar effects of the substituent constants σ‟ are those of Roberts and
Moreland122. Reactions which proceed without significant direct mesomeric
31
interaction between the substituent and the reaction centre, as in the case of
methanolysis of meta- and para- substituted 1-methyl benzoates123 alone
obey the Hammett equation.
CO2 menthyl
R
COOMe
+ CH3O-
+ 1-menthylo-
… (21)
There are several cases in which the Hammett equation fails. In those
cases the chief case for the failure is the direct substituent-reaction centre
interaction. The inductive effect for a substituent, the electron distribution
through σ-bonds, should remain constant. The resonance-polar effect is also
independent of the nature of the reactions. In reactions where the transition
state requires higher electron density for its stabilization, it will be supplied
by resonance interaction with the electron releasing substituents in such case
the parameters of the Hammett equation become deficient124 and the
equation fails. Such failures are encountered in the reactions of the following
type125.
CH2N+(CH3)3
CH2Cl
R
+ N (CH3)3
R
+ Cl-
… (22)
32
Solvolysis of phenyl dimethyl carbinyl chlorides126, 127 is often quoted
as the most attractive example in which the breakdown of the Hammett
equation is clearly seen. There are three essentially equivalent procedures
which deal with such cases of these; the easiest is to determine the constants
from tert-cumyl solvolysis127. The substituent parameters derived in this
manner are called σ+ constants.
Some reaction series show a large deviation with even the most
refined modes of applying the Hammett equation. This may (according to
Shorter104) due to three factors.
i. The complexity of the mechanism throughout the reaction series.
ii. A change in the transition state even if the mechanism is the same
throughout the series.
iii. A change in the rate determining step.
1.9
Cyclanols as Substrate
A variety of oxidants were used during the study of kinetics of
oxidation of cyclic alcohols. In most of the cases orientation was the main
investigation to find out a correlation between the ring size of the cyclanols
and their reactivity. The following table 1.1 results could tell how the
reactivity of cyclanols is independent of their ring size.
33
Table 1.1 Effect of ring size on the oxidation of cyclanols
by various oxidants
Relative rates
Alcohols
Cr (VI)128,129
V (V)130
TI (III)131
Br2132
Cyclopentanol
1.60
0.22
0.027
0.90
Cyclohexanol
1.00
1.00
1.000
1.00
Cycloheptanol
2.00
0.45
0.017
2.70
Cyclooctanol
2.20
0.90
0.047
8.00
Kinetics of oxidation of cyclanols by nitrolic acid has been studied 133.
Rao et al. studied the kinetics of osmium (VIII) catalyzed oxidation of
cyclohexanol by periodate134.
Kinetics and mechanism of osmium tetroxide catalyzed oxidation of
cyclohexanol and methyl cyclohexanol by hexacyano ferrate (III) ions have
been studied by Sing et al.135.
Radhakrishnamurthy and Behera136 studied the kinetics of oxidation
of cyclanols to their respective ketones using cerium (IV) in aqueous acetic
acid and pyridine. The reactivity was as follows.
Cyclooctanol > Cycloheptanol > Cyclopentanol > Cyclohexanol
But jumbled order of the reactivity was observed in the kinetics of
oxidation of cyclanols with potassium hexacyano ferrate (III)137. This order
is not in coincidence with I-Strain theory.
34
In both the cases of study, the authors have not reported any
mechanistic study. But a detailed study on this kinetics was attempted by
later workers138 to analyse the reactivity of cyclanols on the basis of Bayer‟s
strain theory. For example, the reactivity of cyclanols was explained on the
assumption that it reacted in flexible form, rather than in the chair form
because of the greater stability of the latter.
A overall second order rate was followed139 in the oxidation kinetics
when barium manganate was the oxidant. The order of reactivity was found
to be
Cyclopentanol > Cyclohexanol > Cyclooctanol > Cycloheptanol
The greater reactivity of cyclopentanol was explained on the basis of
I-strain theory at the same time the lower rates of cycloheptanol and
cyclooctanol were illustrated due to the probable stabilization of the ring size
by trans annular hydrogen bonding.
Kinetics and mechanism of oxidation of cyclanols by potassium di
tellurato argenate (III)140, potassium hexacyano ferrate (II)141, Ce(SO4)2142
and Thallium (III)143 have been studied.
Spectrophotometric studies144 on the oxidation of cyclohexanol with
Mn3+ in acid medium showed first order dependence on the oxidant at high
concentration, increase in H+ also increased the rate. A suitable mechanism
involving an intermediate complex was proposed.
Sudhakar et al.145 followed the kinetics of oxidation of cyclohexanol
on zinc oxide surface in tert-butyl alcohol in the absence of high percentage.
35
The rate was found to increase linearly, both with the amount of zinc oxide
and carbon tetra chloride used.
The kinetics of oxidation of cyclanols by chloramine-T146, N-Bromo
succinimide147, 1-Chlorobenzotrizole148, N-Chloropiperidone149 have been
studied.
Kabilan et al.149 studied the kinetics of oxidation of cyclanols by
BPCC and PCC. The rate of the reaction is first order with respect to
substrate, oxidant and acid.
Kinetics of oxidation of cyclic alcohols by quinolinium dichromate 150
have been studied in acid medium using DMF as the solvent. The order of
reactivity was,
Cyclooctanol > Cycloheptanol > Cyclopentanol > Cyclohexanol
Gurumurthy et al.151 studied the reactivity of cyclanols towards
quinolinium chlorochromate oxidation. The kinetics of oxidation of few
cyclanols with quinolinium chlorochromate has been investigated in aqueous
acetic acid 70% (v/v) solution in the presence of [H+] ions. The order with
respect to QCC, cyclanol and H+ is one each. The relative reactivity viz.,
Cyclohexanol < Cyclopentanol < Cycloheptanol < Cyclooctanol is explained
on the basis of I-Strain theory.
Oxidation of cyclanols with quinolinium dichromate studied by Sekar
et al.152. Quinolinium dichromate oxidizes cyclanols to corresponding cyclic
36
ketones in 50 % (v/v) acetic acid-water medium. The order of reactivity was
6<7<5<8 i.e., Cyclohexanol < Cycloheptanol < Cyclopentanol <
Cyclooctanol.
Sekar153 studied the structure and reactivity of cyclanols towards
nicotinium dichromate oxidation. The order of reaction has been found to be
one with respect to oxidant and second with respect to substrate. The kinetic
data has obtained, the activation parameters have been calculated and a
plausible mechanism has been proposed. The order of reactivity was
6<7<5<8 on the basis of I-Strain theory.
Structure and reactivity of cyclic alcohols towards pyrazinium
chlorochromate oxidation investigated by Sekar154. The oxidation of cyclic
alcohols with pyrazinium chlorochromate in acid medium using acetic acid
as the solvent results to the formation of the corresponding cyclic ketones.
The order in oxidant, substrate and H+ is one each. From the kinetic data, the
activation
parameters
have
been
calculated
and
a
plausible
mechanism has been proposed. The relative reactivity viz., Cyclohexanol,
Cyclopentanol < Cycloheptanol, Cyclooctanol is explained on the basis of
I-strain theory. The rate of higher ring systems like 7 and 8 is higher in
comparison with lower ring system 5 and 6. It can be concluded that the
order of reactivity for oxidation of cyclic alcohols by pyrazinium
chlorochromate could be rationalized on the basis of change in ring strain
involved.
Chapter - II
SCOPE OF THE WORK
The development of new Chromium (VI) reagents for the oxidation of
several organic substrates continues to be a subject of interest.
Chromium (VI) is highly toxic in nature. But, chromium (VI) and
chromium (III) are non-toxic and ecofriendly in nature. Chromium (VI)
compounds act as irritants of the skin and mucous membranes.
Chromium (VI) originating from tanneries and industrial waste contaminates
soil, thus presenting a serious environmental hazard due to its toxicity.
The present work aims at conversion of highly toxic chromium (VI)
into a non-toxic chromium (IV) or chromium (III) compound.
In recent years, kinetics and mechanism of oxidation reactions of
chromium (VI) for a number of substrates have been fairly well studied. The
interesting point in the chromium (VI) oxidation is the mechanism of
oxidation varies with the nature of the chromium (VI) species and the
solvent used. Quinoxalinium dichromate is one among the oxidant and due
to its synthetic and selective oxidant nature. It has been proposed to
investigate the kinetic and mechanistic aspects of oxidation reaction of
quinoxalinium dichromate.
Study of the oxidation of S-phenylmercaptoacetic acid by various
oxidants is reported in literature. There seems to be no report on the
oxidation of S-phenylmercaptoacetic acid by quinoxalinium dichromate. It is
proposed to investigate the kinetics and mechanism of oxidation of
S-phenylmercaptoacetic acid in presence of perchloric acid with a view to
38
understand about the substituent effect and possible mechanism for the
reaction.
A literature survey revealed that a considerable attention has been
shown on the kinetics and mechanism of oxidation of cyclanols by various
oxidants. There has been no report on the quinoxalinium dichromate of
cyclanols. Hence, it is proposed to study the oxidation of cyclanols in
presence of perchloric acid with view to understand about the reactivity and
a suitable mechanism proposed for the reaction.
2.1
I-Strain Effect in Ring Compounds
The six membered ring derivative cyclohexanol is highly reactive
towards carbonyl reagents, cyanide ion and semicarbazide. Whereas
cyclopentanol and cyclooctanol are surprisingly inert towards these reagents.
The difference in reactivity between 5- and 7- membered ring compounds on
the one hand and the 6- membered ring compounds on the other are quite
large and involve factors of the order of 100. The difference in reactivity
points to an important effect of ring size on chemical behavior i.e., I-strain.
I-strain is that change in internal strain which results from the change in
coordination number of ring atoms involved in a chemical reaction.
Chapter - III
EXPERIMENTAL METHODS
.
3.1
Preparation of Quinoxalinium Dichromate
A solution of 26.4 g quinoxaline (0.2 mol) in 60 ml water was slowly
added to a cooled solution of 21.0 g chromium trioxide (0.2 mol) in 20 ml
water. After 30 min, the reaction mixture was diluted with 40 ml acetone,
cooled to -15 oC, and the orange solid was filtered, washed with acetone, and
dried in vacuo. The compound melted at (116 ºC) (lit155 m.p. 115-116 ºC);
yield: 78% and further analysed through spectral studies.
3.2
Preparation of S-Phenylmercaptoacetic Acid
A procedure similar to that of Gabriel156 was adopted to prepare the
substituted S-phenylmercaptoacetic acids.
Chloroacetic acid (5 g) was dissolved in a solution of sodium
hydroxide (5.5 g in 30 ml of water). To the resulting solution was added the
thiophenol (0.05 mol) without allowing the temperature to rise. The mixture
was heated under reflux in an oil bath for 3-4 h, allowed to cool and
acidified with concentrated hydrochloric acid. Oil separated and this
solidified on cooling.
The solid was filtered off and dissolved in a dilute solution of sodium
carbonate. The alkali layer was shaken with ether to remove the unreacted
thiophenol. The aqueous solution was neutralized with concentrated
hydrochloric acid and the separated acid was recrystallized from Benzene
petroleum ether. The compound melted at 63 ºC (lit157 m.p. 63.5 - 64 ºC).
40
The substituted S-phenylmercaptoacetic acids were prepared by the
following procedures.
3.2.1 Preparation of meta- methoxy S-Phenylmercaptoacetic Acid
Chloroacetic acid (5 g) was dissolved in a solution of sodium
hydroxide (5.5 g in 30 ml of water). To the resulting solution was added the
meta- methoxy thiophenol (0.05 mol) without allowing the temperature to
rise. The mixture was heated under reflux in an oil bath for 3-4 h, allowed to
cool and acidified with concentrated hydrochloric acid. Oil separated and
this solidified on cooling.
The solid was filtered off and dissolved in a dilute solution of sodium
carbonate. The alkali layer was shaken with ether to remove the unreacted
thiophenol. The aqueous solution was neutralized with concentrated
hydrochloric acid and the separated acid was recrystallized from acetic acidwater. The compound melted at 62 ºC (lit158 m.p. 62.5 ºC).
3.2.2 Preparation of para- methoxy S-Phenylmercaptoacetic Acid
Chloroacetic acid (5 g) was dissolved in a solution of sodium
hydroxide (5.5 g in 30 ml of water). To the resulting solution was added the
para- methoxy thiophenol (0.05 mol) without allowing the temperature to
rise. The mixture was heated under reflux in an oil bath for 3-4 h, allowed to
cool and acidified with concentrated hydrochloric acid. Oil separated and
this solidified on cooling.
The solid was filtered off and dissolved in a dilute solution of sodium
carbonate. The alkali layer was shaken with ether to remove the unreacted
41
thiophenol. The aqueous solution was neutralized with concentrated
hydrochloric acid and the separated acid was recrystallized from acetic acidwater. The compound melted at 76 ºC (lit157 m.p. 75.8-76.2 ºC).
3.2.3 Preparation of meta- methyl S-Phenylmercaptoacetic Acid
Chloroacetic acid (5 g) was dissolved in a solution of sodium
hydroxide (5.5 g in 30 ml of water). To the resulting solution was added the
meta- methyl thiophenol (0.05 mol) without allowing the temperature to rise.
The mixture was heated under reflux in an oil bath for 3-4 h, allowed to cool
and acidified with concentrated hydrochloric acid. Oil separated and this
solidified on cooling.
The solid was filtered off and dissolved in a dilute solution of sodium
carbonate. The alkali layer was shaken with ether to remove the unreacted
thiophenol. The aqueous solution was neutralized with concentrated
hydrochloric acid and the separated acid was recrystallized from acetic acidwater. The compound melted at 66 ºC (lit157 m.p. 66.8-67.4 ºC).
3.2.4 Preparation of para- methyl S-Phenylmercaptoacetic Acid
Chloroacetic acid (5 g) was dissolved in a solution of sodium
hydroxide (5.5 g in 30 ml of water). To the resulting solution was added the
para- methyl thiophenol (0.05 mol) without allowing the temperature to rise.
The mixture was heated under reflux in an oil bath for 3-4 h, allowed to cool
and acidified with concentrated hydrochloric acid. Oil separated and this
solidified on cooling.
The solid was filtered off and dissolved in a dilute solution of sodium
carbonate. The alkali layer was shaken with ether to remove the unreacted
42
thiophenol. The aqueous solution was neutralized with concentrated
hydrochloric acid and the separated acid was recrystallized from ethanol.
The compound melted at 94 ºC (lit157 m.p. 94-94.4 ºC).
3.2.5 Preparation of meta- bromo S-Phenylmercaptoacetic Acid
Chloroacetic acid (5 g) was dissolved in a solution of sodium
hydroxide (5.5 g in 30 ml of water). To the resulting solution was added the
meta- bromo thiophenol (0.05 mol) without allowing the temperature to rise.
The mixture was heated under reflux in an oil bath for 3-4 h, allowed to cool
and acidified with concentrated hydrochloric acid. Oil separated and this
solidified on cooling.
The solid was filtered off and dissolved in a dilute solution of sodium
carbonate. The alkali layer was shaken with ether to remove the unreacted
thiophenol. The aqueous solution was neutralized with concentrated
hydrochloric acid and the separated acid was recrystallized from acetic acidwater. The compound melted at 85 ºC (lit159 m.p. 85-86 ºC).
3.2.6 Preparation of para- bromo S-Phenylmercaptoacetic Acid
Chloroacetic acid (5 g) was dissolved in a solution of sodium
hydroxide (5.5 g in 30 ml of water). To the resulting solution was added the
para- bromo thiophenol (0.05 mol) without allowing the temperature to rise.
The mixture was heated under reflux in an oil bath for 3-4 h, allowed to cool
and acidified with concentrated hydrochloric acid. Oil separated and this
solidified on cooling.
The solid was filtered off and dissolved in a dilute solution of sodium
carbonate. The alkali layer was shaken with ether to remove the unreacted
43
thiophenol. The aqueous solution was neutralized with concentrated
hydrochloric acid and the separated acid was recrystallized from water. The
compound melted at 117 ºC (lit157 m.p. 118-118.5 ºC).
3.2.7 Preparation of meta- chloro S-Phenylmercaptoacetic Acid
Chloroacetic acid (5 g) was dissolved in a solution of sodium
hydroxide (5.5 g in 30 ml of water). To the resulting solution was added the
meta- chloro thiophenol (0.05 mol) without allowing the temperature to rise.
The mixture was heated under reflux in an oil bath for 3-4 h, allowed to cool
and acidified with concentrated hydrochloric acid. Oil separated and this
solidified on cooling.
The solid was filtered off and dissolved in a dilute solution of sodium
carbonate. The alkali layer was shaken with ether to remove the unreacted
thiophenol. The aqueous solution was neutralized with concentrated
hydrochloric acid and the separated acid was recrystallized from acetic acidwater. The compound melted at 81 ºC (lit157 m.p. 81.5-82.2 ºC).
3.2.8 Preparation of para- chloro S-Phenylmercaptoacetic Acid
Chloroacetic acid (5 g) was dissolved in a solution of sodium
hydroxide (5.5 g in 30 ml of water). To the resulting solution was added the
para- chloro thiophenol (0.05 mol) without allowing the temperature to rise.
The mixture was heated under reflux in an oil bath for 3-4 h, allowed to cool
and acidified with concentrated hydrochloric acid. Oil separated and this
solidified on cooling.
The solid was filtered off and dissolved in a dilute solution of sodium
carbonate. The alkali layer was shaken with ether to remove the unreacted
44
thiophenol. The aqueous solution was neutralized with concentrated
hydrochloric acid and the separated acid was recrystallized from water. The
compound melted at 106 ºC (lit157 m.p. 105-107 ºC).
3.2.9 Preparation of meta- nitro S-Phenylmercaptoacetic Acid
Chloroacetic acid (5 g) was dissolved in a solution of sodium
hydroxide (5.5 g in 30 ml of water). To the resulting solution was added the
meta- nitro thiophenol (0.05 mol) without allowing the temperature to rise.
The mixture was heated under reflux in an oil bath for 3-4 h, allowed to cool
and acidified with concentrated hydrochloric acid. Oil separated and this
solidified on cooling.
The solid was filtered off and dissolved in a dilute solution of sodium
carbonate. The alkali layer was shaken with ether to remove the unreacted
thiophenol. The aqueous solution was neutralized with concentrated
hydrochloric acid and the separated acid was recrystallized from acetic acidwater. The compound melted at 136 ºC (m.p. 136-137 ºC).
3.2.10 Preparation of para- nitro S-Phenylmercaptoacetic Acid
Chloroacetic acid (5 g) was dissolved in a solution of sodium
hydroxide (5.5 g in 30 ml of water). To the resulting solution was added the
para- nitro thiophenol (0.05 mol) without allowing the temperature to rise.
The mixture was heated under reflux in an oil bath for 3-4 h, allowed to cool
and acidified with concentrated hydrochloric acid. Oil separated and this
solidified on cooling.
The solid was filtered off and dissolved in a dilute solution of sodium
carbonate. The alkali layer was shaken with ether to remove the unreacted
45
thiophenol. The aqueous solution was neutralized with concentrated
hydrochloric acid and the separated acid was recrystallized from ethanolwater.
The compound melted at 152 ºC (lit160 m.p. 151-153 ºC). The
physical constants of some substituted S-phenylmercaptoacetic acids are
listed in table 1.2.
Table 1.2 Physical constants of substituted S-phenylmercaptoacetic acid
Compound
Molecular Formula
Melting Point 0C
Observed
Literature
C6H5SCH2COOH
63
63.5-64
m-OCH3C6H4SCH2COOH
62
62.5
p-methoxy phenylmercaptoacetic acid
p-OCH3C6H4SCH2COOH
76
75.8-76.2
m-methyl phenylmercaptoacetic acid
m-CH3C6H4SCH2COOH
66
66.8-67.4
p-methyl phenylmercaptoacetic acid
p-CH3C6H4SCH2COOH
94
94-94.4
m-bromo phenylmercaptoacetic acid
m-BrC6H4SCH2COOH
85
85-86
p-bromo phenylmercaptoacetic acid
p-BrC6H4SCH2COOH
117
118-118.5
m-chloro phenylmercaptoacetic acid
m-ClC6H4SCH2COOH
81
81.5-82.2
p-chloro phenylmercaptoacetic acid
p-ClC6H4SCH2COOH
106
105-107
m-nitro phenylmercaptoacetic acid
m-NO2C6H4SCH2COOH
136
136-137
p-nitro phenylmercaptoacetic acid
p-NO2C6H4SCH2COOH
152
151-153
S-phenylmercaptoacetic acid
m-methoxy phenylmercaptoacetic acid
46
3.3
Determination of Physical Constants of Cyclanols
The following Cyclanols (Aldrich, Fluka grade) were used as such
and their Physical constants are given in Table 1.3
Table 1.3 Physical constants of Cyclanols
Boiling Point 0C
Compound
Molecular Formula
Melting Point 0C
Observed
Literature
Observed
Literature
Cyclohexanol
C6H12O
160
161.84
---
---
Cyclopentanol
C5H10O
140
139-140
---
---
Cycloheptanol
C7H14O
184
185
---
---
Cyclooctanol
C8H16O
---
---
107
106-108
Purification of Acetic acid
The procedure followed for the purification of acetic acid was
essentially similar to that of Weissberger161. Two litres of glacial acetic acid
(AR) was partially frozen and about one litre of the liquid was removed.
The residue was melted and refluxed with chromium trioxide (30 g) for
4 h and fractionally distilled. The portion distilling between 116-118 ºC was
collected, partially frozen and about half of the acid was discarded as liquid.
The remaining residue was melted and fractionated again after treating with
chromium trioxide (30 g). The fraction boiling at 116-118 ºC was collected
and kept in brown bottles.
47
Double Distilled Water
Deionised water was distilled twice in „corning‟ glass vessels, the
second distillation being from alkaline potassium permanganate and was
used throughout the kinetic measurements.
Other Reagents
Perchloric acid, sodium perchlorate, acrylonitrile, manganous
sulphate, sodium thiosulphate, potassium iodide and starch were all of
AnalaR grade (E-merck) and were used as such.
3.4
Kinetic
Measurements
for
the
Oxidation
S-Phenylmercaptoacetic Acid by Quinoxalinium Dichromate
of
Solutions of S-phenylmercaptoacetic acid in acetic acid and other
reagent like quinoxalinium dichromate, perchloric acid solutions in doubly
distilled water were prepared. In all the reactions pseudo-first order
conditions were maintained. The kinetic measurements were made using
spectrophotometer (ELICO SL 207 MINI SPEC) λmax = 470 nm as follows.
All the solutions were kept in a thermostat constant temperature for
half an hour for each run. The temperature was controlled using Raagaa
thermostat to an accuracy of ± 0.1˚C. Then the reaction was started by
adding a known volume of quinoxalinium dichromate into the reaction flask.
Immediately,
1 ml of aliquot (approximate) was transferred in to the quartz
cuvette, which had already been kept and thermostated in the instrument.
The reactions were followed by determining the concentration of the
unreacted quinoxalinium dichromate for known intervals of time.
48
Evaluation of Rate Constants
The reactions were carried out under pseudo-first order conditions,
keeping the substrate concentration always in excess. The pseudo-first order
rate constant of each kinetic run was evaluated from the slope of the linear
plot of log absorbance versus time, according to the first order rate equation
by the method of least squares.
t = 2.303 / k1 × log (a/a-x)
k1 = 2.303 / t × log (a/a-x)
k1 = 2.303 × slope, expressed in s-1 and „a‟ and (a-x) denote the initial
concentration and the concentration at time „t‟ respectively of quinoxalinium
dichromate. The linearity of each fit is given in terms of the correlation
co-efficient (r).
Evaluation of Thermodynamic Parameters
The enthalpy (ΔH#) and entropy of activation (ΔS#) of a reaction are
related to the specific reaction rate (k′) and absolute temperature (T) by the
Eyring‟s equation:
kBT e-ΔH#
RT
e ΔS#
R
k‟ =
h
where, kB = the Boltzmann constant; h = the Planck‟s constant.
These activation parameters were calculated by the least square
analysis of a linear plot of ln (k/T) versus (1/T) of the Eyring‟s equation ΔH#
and ΔS# were calculated from the slope and intercept of the plot
respectively, as per the following equation:
49
ΔS#
ln k‟/ T = 23.7604 +
0.008314
ΔH#
0.008314
1
T
ΔH# = 0.008314 × (slope) kJ mol-1
ΔS# = 0.008314 × (intercept-23.7604) JK-1 mol-1
The free energy of activation and energy of activation were calculated using
the thermodynamic relationship.
ΔG# = ΔH# -TΔS kJ mol-1
Ea = ΔH# + RT kJ mol-1
Accuracy of the results
The pseudo - first order rate constants were calculated by the method
of least squares. Duplicate runs were carried out and the results were found
to be reproducible within ±3%.
3.5
Non-Kinetic Study
Stoichiometry
The kinetics of reaction was to establish the stoichiometry of the
reaction and identify any side reactions. The stoichiometry of the reaction
[QxDC]:[S-phenylmercaptoacetic aicd] was determined by taking excess of
[QxDC] over [S-phenylmercaptoacetic aicd] and allowing the reaction to go
for completion. After sufficient length of time, all the substrate has
completely reacted to quinoxalinium dichromate leaving behind the
unreacted
quinoxalinium dichromate.
The
unreacted
quinoxalinium
dichromate was estimated iodometrically. The estimation of unreacted
50
quinoxalinium dichromate showed that one mole of substrate consumed by
one mole of oxidant. Thestoichiometry between S-phenylmercaptoacetic
acid and QxDC was found to be 1:1.
Product analysis
The reaction mixture containing S-phenylmercaptoacetic aicd (0.1 M)
in acetic acid and QxDC (0.1 M) in acetic acid was added and the medium
was maintained using perchloric acid. Then the reaction mixture was slightly
warmed and was kept aside for about 48 h for the completion of reaction.
After 48 h, the reaction mixture was extracted with ether and dried over
anhydrous sodium sulphate. The ether layer was washed with water several
times and kept on a water bath for ether evaporation and cooled to get the
residue.
The residue was subjected to TLC analysis on a silica gel plate
developed in a solvent system of n-butanol-acetic acid –water (40 to 50%,
upper layer was used). The residue gave two spots, which were made visible
by exposure to iodine; one corresponding to (phenylmercapto)acetic acid
(Rf = 0.84) and the other to phenylsulphinylacetic acid (Rf = 0.45). Further,
the IR Spectra of the residue showed an intense absorption band at
1030 cm-1 characteristic acid of =S=O, Stretching frequency.
51
3.6
Kinetic Measurements for the Oxidation of Cyclanols by
Quinoxalinium Dichromate
Solution of cyclohexanol in acetic acid and other reagents like
quinoxalinium dichromate, sodium perchlorate and perchloric acid solutions
were prepared in doubly distilled water. In all the reactions pseudo - first
order conditions were maintained. The kinetic measurements were made
using
spectrophotometer (ELICO SL 207 MINI SPEC) λmax = 470 nm as
follows.
Then the reaction was started by adding a known volume of
quinoxalinium dichromate into the reaction flask. Immediately, 1ml of
aliquot (approximate) was transferred into the quartz cuvette which had
already been kept and thermostated in the instrument. The reactions were
followed by determining the concentration of the unreacted quinoxalinium
dichromate, for known intervals of time. The pseudo - first order rate
constants were obtained from the slopes of the log absorbance versus time
plots. (for each kinetic run number of data points = 12; correlation
coefficient = 0.999).
3.7
Non-Kinetic Study
Stoichiometry
The kinetics of reaction was to establish the stoichiometry of the
reaction and identify any side reactions. The stoichiometry of the reaction
QxDC:cyclanols was determined by taking excess of QxDC over cyclanols
and allowing the reaction to go for completion. After sufficient length of
time, all the substrate has completely reacted to quinoxalinium dichromate
leaving behind the unreacted quinoxalinium dichromate. The unreacted
52
quinoxalinium dichromate was estimated iodometrically. The estimation of
unreacted quinoxalinium dichromate showed that one mole of substrate
consumed by one mole of oxidant. The stoichiometry between cyclanols and
QxDC was found to be 1:1.
Product analysis
The reaction mixture containing cyclohexanol (0.1 M) in acetic acid
and QxDC (0.1 M) in acetic acid was added and the medium was maintained
using perchloric acid. Then the reaction mixture was slightly warmed and
was kept aside for about 48 h for the completion of reaction. After 48 h, the
reaction mixture was extracted with ether and dried over anhydrous sodium
sulphate. The ether layer was washed with water several times and kept on a
water bath for ether evaporation and cooled to get the product. The product
was identified as cyclohexanone by its IR and mass spectral studies (Plate 1
& Plate 2).
53
Plate.1 IR Spectrum of Cyclohexanone
54
Plate.2 Mass Spectrum of Cyclohexanone
Chapter - IV
RESULTS AND DISCUSSION
..
4.1
Kinetics
and
Mechanism
of
Oxidation
of
S-Phenylmercaptoacetic Acid by Quinoxalinium Dichromate in
Aqueous Acetic Acid Medium
The kinetics of oxidation of S-phenylmercaptoacetic acid with
quinoxalinium dichromate in protic solvent system was carried out under
psuedo-first order conditions. The results are discussed in the following
pages.
Effect of varying the [QxDC]
The reaction was investigated with varying concentrations of
quinoxalinium dichromate at constant S-phenylmercaptoacetic acid and
perchloric acid concentrations. The reaction was found to be first order with
respect to QxDC as evidenced by the linear plot of log absorbance versus
time (Fig 4.1) and also from the constancy of the pseudo-first order rate
constants. The rate or reaction decreased with increase the concentration of
quinoxalinium
dichromate.
Because
in
that
condition
the
total
chromium (VI) was in the form of acid chromate in which is the effective
oxidant162.
56
Table 4.1
[PMA] = 5.00×10-2mol dm-3
AcOH-H2O = 50:50 (%)
[H+] = 3.50×10-1mol dm-3
Temperature = 313 K
[QxDC] 103
mol dm-3
1.5
k1 104
s-1
3.99
2.0
3.26
2.5
2.35
3.0
1.97
3.5
1.62
57
Fig 4.1 Plot of log absorbance versus time
58
Effect of varying the [PMA]
At a constant temperature, the rate increased steadily on increasing the
concentration of the substrate as shown in Table 4.2. The linear plot of log k
versus log [substrate] with a slope of unity (Fig.4.2) clearly indicates that the
reaction has unit order dependence on the concentration of the substrate. The
specific reaction rate constant of k2 = k1/[s] confirms the first order in the
S-phenylmercaptoacetic acid.
Table 4.2
[QxDC] = 2.00×10-3mol dm-3
AcOH-H2O = 50:50 (%)
[H+]
Temperature = 313 K
= 3.50×10-1mol dm-3
[PMA] 102
mol dm-3
2.5
k1 104
s-1
1.27
k2 = k1/[s] 10
mol-1 dm 3 s-1
0.05
5.0
3.26
0.06
7.5
4.42
0.06
10.0
6.24
0.06
12.5
8.52
0.07
15.0
10.38
0.07
59
r=0.995
B=1.14
Fig 4.2 Plot of log k versus log [substrate]
60
Effect of varying hydrogen ion concentration
The kinetic runs were performed at different concentrations of
perchloric acid which acted as the catalyst. The rate of reaction decreased
with increase the concentration of hydrogen ion. Plots of k versus 1/[H+] and
log k versus log [H+] (Fig. 4.3 and Fig. 4.4) are also straight lines with unit
slope indicating an inverse first order dependence on hydrogen ion
concentration163.
Table 4.3
[QxDC] = 2.00×10-3mol dm-3
AcOH-H2O = 50:50 (%)
[PMA] = 5.00×10-2mol dm-3
Temperature = 313 K
[HClO4] 101
mol dm-3
3.5
k1 104
s-1
3.26
7.0
1.98
10.5
1.46
14.0
0.98
17.5
0.78
61
Fig 4.3 Plot of k versus 1 / [H+]
62
Fig 4.4 Plot of log k versus log [H+]
63
Effect of varying the ionic strength
The reaction was carried out at different initial concentrations of
sodium perchlorate while the other variables were kept constant. Increase in
ionic strength of the medium by adding sodium perchlorate has no effect on
the reaction rate indicating the participation of charged species as a reactant
in the rate - determining step.
Table 4.4
[QxDC] = 2.00×10-3mol dm-3
AcOH-H2O = 50:50 (%)
[PMA] = 5.00×10-2mol dm-3
Temperature = 313 K
[H+]
=3.50×10-1mol dm-3
[NaClO4] 102
mol dm-3
0.00
k1 104
s-1
3.26
5.05
3.24
10.10
3.18
15.15
3.22
20.20
3.16
64
Effect of varying solvent composition
The acetic acid composition in the solvent mixture was varied while
maintaining the other variables constant. The rate was found to increase
considerably on increasing the acetic acid content of the medium. The plot
of log k1 versus 1/D (Fig. 4.5) gave a straight line with a positive slope154.
This might be probably due to ion-dipole interaction in the rate determining
step.
Table 4.5
[QxDC] = 2.00×10-3mol dm-3
[H+] = 3.50×10-1mol dm-3
[PMA] = 5.00×10-2mol dm-3
Temperature = 313 K
AcOH-H2O
% (v/v)
40-60
D
k1 104
s-1
49.60
2.87
45-55
45.99
2.99
50-50
42.37
3.26
55-45
38.75
3.48
60-40
35.14
3.88
65
Fig 4.5 Plot of log k versus 1 / D
66
Effect of added acrylonitrile
The addition of acrylonitrile, which is a very good trapper of free
radicals, did not have any retarding effect on the reaction. It indicates that no
free radicals participation in the reaction.
Effect of varying the [Manganous Sulphate]
The kinetic runs were carried out at different concentration of
manganous sulphate and maintain the concentration of substrate, oxidant,
hydrogen ion, dielectric constant and temperature of the reaction. The rate of
reaction decreases with increase in the concentration of manganous sulphate.
Thus, it is possible that the reaction involves a two-electron transfer process
in the mechanism.
Table 4.6
[QxDC] = 2.00×10-3mol dm-3
AcOH-H2O = 50:50 (%)
[PMA] = 5.00×10-2mol dm-3
Temperature = 313 K
[H+]
= 3.50×10-1mol dm-3
[MnSO4] 102
mol dm-3
0.00
k1 104
s-1
3.26
0.15
1.99
0.30
1.64
0.45
1.18
0.60
0.97
67
Effect of varying the temperature
The reaction has been carried out at four different temperatures
keeping all the other factors constant. The thermodynamic parameters have
been computed163 from the linear plot of log (k2/T) versus 1/T of Eyring‟s
equation164 (Fig 4.6)
ΔH#= 39.30 kJ mol-1
ΔS#= -125.06 JK-1mol-1
ΔG# = 78.44 kJ mol-1at 313 K
Table 4.7
[QxDC] = 2.00×10-3mol dm-3
[H+] = 3.50×10-1mol dm-3
[PMA] = 5.00×10-2mol dm-3
AcOH-H2O = 50:50 (%)
Temperature
K
303
k1 104
s-1
2.54
313
3.26
323
5.58
333
8.21
68
r = 0.989
Fig 4.6 Plot of log k2/T versus 1/T
69
Mechanism
The oxidation of S-phenylmercaptoacetic acid with quinoxalinium
dichromate was catalyzed by perchloric acid. It is first order with respect to
the concentrations of each of the oxidant and substrate and inverse first order
with respect to H+. Product analysis clearly indicates that the obtaining of
the corresponding sulphoxide. From these observations, the following
mechanism and rate law were proposed. (Scheme 4.1).
K1
+
(C8H7N2)Cr2O7H
K2
+
(C8H7N2)Cr2O7H + C6H5SCH2COOH
k3
complex
(C8H7N2)Cr2O7 + H+
complex
O
C6H5– S –CH2COOH + H+ + Cr (IV)
Scheme 4.1
Rate law:
Rate = k3 [complex]
= k3K2 [QxDCH+] [PMA]
= k3K2K1 [QxDC] [PMA] [H+] / K2 [H+]
-d[QxDC]/dt = K1K2k3 [QxDC] [PMA] [H+] / K2 [H+]
70
The proposed mechanism and the rate law support all the observations
made including the effect of solvent polarity and the negative entropy of
activation.
4.2.
Oxidation of Substituted
Quinoxalinium Dichromate
S-Phenylmercaptoacetic
Acid
by
The kinetics of oxidation of substituted S-phenylmercaptoacetic acid
with substituents such as -methoxy, -methyl, -bromo, -chloro and -nitro were
carried out employing different initial concentrations of the substituted
S-phenylmercaptoacetic acid. The pseudo-first order rate constant obtained
for those substrates showed unit order dependence with respect to
substituted S-phenylmercaptoacetic acid. The effect of varying substrate
concentration on rate constants and varying temperature are shown in
Table 4.8.
Table 4.8. Effect of varying the concentration of substrate and
temperature
[QxDC] = 2.00×10-3mol dm-3
[H+]
AcOH-H2O = 50:50 (%)
= 3.50×10-1mol dm-3
Substituent Temperature [substrate]102 k1 104
K
mol dm-3
s-1
p-OCH3a
313
2.5
0.62
313
5.0
1.12
313
7.5
1.65
313
10.0
2.04
313
12.5
2.55
313
15.0
3.06
Order with
respect
to substrate
0.88
2
k1 104
Substituent Temperature [substrate]10
K
s-1
mol dm-3
p-OCH3b
p-CH3a
p-CH3b
p-Bra
p-Brb
303
5.0
0.72
313
5.0
1.12
323
5.0
1.85
333
5.0
2.71
313
2.5
0.82
313
5.0
1.55
313
7.5
2.17
313
10.0
2.88
313
12.5
3.54
313
15.0
4.18
303
5.0
0.94
313
5.0
1.55
323
5.0
2.38
333
5.0
3.56
313
2.5
3.96
313
5.0
9.43
313
7.5
13.39
313
10.0
18.37
313
12.5
25.03
313
15.0
31.08
303
5.0
5.98
313
5.0
9.43
323
5.0
13.89
333
5.0
19.98
Order with
respect
to substrate
-
0.90
-
1.11
-
Substituent Temperature [substrate]10
K
mol dm-3
p-Cla
p-Clb
p-NO2a
p-NO2b
m-OCH3a
2
k1 104
s-1
313
2.5
3.08
313
5.0
6.76
313
7.5
10.89
313
10.0
14.99
313
12.5
19.27
313
15.0
23.04
303
5.0
5.04
313
5.0
6.76
323
5.0
8.87
333
5.0
10.98
313
2.5
25.01
313
5.0
45.71
313
7.5
68.55
313
10.0
86.89
313
12.5
98.40
313
15.0
108.36
303
5.0
30.12
313
5.0
45.71
323
5.0
64.94
333
5.0
87.22
313
2.5
2.97
313
5.0
4.84
313
7.5
6.52
313
10.0
8.24
313
12.5
10.17
313
15.0
11.96
Order with
respect
to substrate
1.14
0.87
-
0.79
Substituent Temperature [substrate]10
K
mol dm-3
m-OCH3b
m-CH3a
m-CH3b
m-Bra
m-Br
2
k1 10
s-1
4
303
5.0
3.01
313
5.0
4.84
323
5.0
6.52
333
5.0
9.27
313
2.5
1.91
313
5.0
2.98
313
7.5
4.10
313
10.0
5.72
313
12.5
7.14
313
15.0
8.66
303
5.0
1.89
313
5.0
2.98
323
5.0
5.81
333
5.0
8.94
313
2.5
6.74
313
5.0
14.45
313
7.5
22.28
313
10.0
30.13
313
12.5
39.08
313
15.0
48.14
303
5.0
10.21
313
5.0
14.45
323
5.0
19.21
333
5.0
25.09
b
Order with
respect
to substrate
-
0.82
-
1.08
-
Temperature [substrate]102 k1 104
Substituent
K
s-1
mol dm-3
m-Cla
m-Clb
m-NO2a
m-NO2b
313
2.5
5.61
313
5.0
12.88
313
7.5
20.28
313
10.0
28.51
313
12.5
38.72
313
15.0
49.32
303
5.0
9.07
313
5.0
12.88
323
5.0
16.62
333
5.0
21.08
313
2.5
16.35
313
5.0
38.90
313
7.5
63.79
313
10.0
92.31
313
12.5
133.22
313
15.0
174.18
303
5.0
28.21
313
5.0
38.90
323
5.0
50.09
333
5.0
65.24
Order with
respect
to substrate
1.18
-
1.28
-
75
From the above table 4.8, it is clear that among the substituents
studied show a unit order dependence on the reaction rate. From a series of
kinetic runs the rate constants (k1) were estimated for the above substituted
S-phenylmercaptoacetic acid at four different temperatures viz., 303 K,
313 K, 323 K and 333 K in order to calculate the thermodynamic parameters
given in Table 4.9.
76
Table 4.9 Thermodynamic parameters for the oxidation of para- and meta- substituted
S-phenylmercaptoacetic acids by QxDC
[PMA] = 5.00×10-2mol dm-3
[H+] = 3.50×10-1mol dm-3
[QxDC] = 2.00 × 10-3mol dm-3
AcOH-H2O = 50-50 (%)
k1 104
(s-1)
-ΔS#
(JK-1mol-1)
ΔG#
(kJ mol-1)
at 313 K
r
13.60
184.32
71.29
0.990
2.71
15.19
183.41
72.60
0.999
2.38
3.56
15.02
182.91
72.27
0.999
9.43
13.89
19.98
13.48
181.33
70.24
0.999
5.04
6.76
8.87
10.98
8.37
196.29
69.81
0.998
0.87
30.12
45.71
64.94
87.22
11.79
180.99
68.44
0.998
m-OCH3
0.79
3.01
4.84
6.52
9.27
12.37
187.40
71.03
0.996
8
m-CH3
0.82
1.89
2.98
5.81
8.94
18.28
169.86
71.45
0.996
9
m-Br
1.08
10.21
14.45
19.21
25.09
9.73
191.72
69.74
0.999
10
m-Cl
1.18
9.07
12.88
16.62
21.08
9.03
194.47
69.90
0.997
11
m-NO2
1.28
28.21
38.90
50.09
65.24
8.95
190.64
68.62
0.999
Order with
respect to
substrate
303 K
313 K
323 K
333 K
ΔH
(kJ mol-1)
S.
No.
Substituents
1
-H
1.14
2.54
3.26
5.58
8.21
2
p-OCH3
0.88
0.72
1.12
1.85
3
p-CH3
0.90
0.94
1.55
4
p-Br
1.11
5.98
5
p-Cl
1.14
6
p-NO2
7
#
77
The negative values of the entropies of activation (ΔS#) suggested that
the transition state formed was considerable rigid, resulting in a reduction in
the degree of freedom of the molecules. The constancy of the (ΔG#) values
indicated a common mechanism for the oxidation of all the substrates. As
ΔH# and ΔS# do not vary linearly, no isokinetic relationship is observed.
This indicated the absence of enthalpy - entropy compensation effect.
Exner165-167 criticized the validity of such a linear correlation between
ΔH# and ΔS#, as these quantities are dependent on each other. When
measurements at two different temperatures have been made, the data can be
analysed by the following equation168-170.
log (k1)T2 = a + b log (k1)T1
where, a and b are intercept and slope and T2 > T1
The plot of log k313 K versus log k303 K (Fig.4.7) gave a straight line
with an excellent correlation co efficient r = 0.997. Such a good correlation
indicates that all the S-phenylmercaptoacetic acid follow a common
mechanism. The low Ea and ΔH# values support the proposed concerted
mechanism. The negative values of the entropy of activation (ΔS#) suggest
assumption of highly solvated transition state due to its increased polarity.
The reactions are characterised by near constancy in ΔG# values and
excellent linearity in the Exner‟s plot as well as isokinetic plot supporting
the operation of a similar mechanism in the reaction series.
78
Fig 4.7 Exner plot of log k(313 K) versus log k(303 K) for the oxidation of
S-phenylmercaptoacetic acids by quinoxalinium dichromate
(Numbers as given in Table 4.9)
79
4.3
Correlation and reactivity - Effect of substituents
Substituent effects are used to probe into reaction mechanism since
the aromatic system can be affected by electronic nature of the substituents.
The effect of substituents and the reactivity has been investigated by
employing ten meta- and para- substituted S-phenylmercaptoacetic acid.
From the results it is evident that electron withdrawing substituents are
found to enhance the reaction rate and electron releasing substituents are
found to retard the rate of reaction171-176. The rate constant for all the
substituents and activation parameters were evaluated from the Eyring‟s
plots and are listed in Table 4.9.
The correlation of log kobs with Hammett‟s substituent constant177, 178
σ gave a linear plot of a positive slope179, 180 (r = 0.994,  = +1.54) (Fig 4.8)
with the Hammett value at 313 K. The positive „ρ‟ value indicates that
electron-withdrawing substituents enhance the rate of oxidation and
electron-releasing substituents decrease the rate of the reaction. Correlation
of rate data with σ is also satisfactory with r = 0.994, suggesting a reaction
centre has higher electron density in the transition state than in the starting
material.
80
Fig 4.8 Plot of log k(313K) against Hammett‟s substituent
constant σ in the Oxidation of S-phenylmercaptoacetic acid by QxDC
(Numbers as given in table 4.9)
81
4.4.
Kinetics and Mechanism of Oxidation of Cyclohexanol by
Quinoxalinium Dichromate in Aqueous Acetic acid Medium
In order to obtain a clear picture of the mechanism of oxidation of
cyclohexanol by quinoxalinium dichromate in aqueous acetic acid solvent
system the reaction was carried out in pseudo - first order conditions. The
results are discussed in the following pages.
Effect of varying the [QxDC]
At fixed [H+] with [substrate] in excess, the plot of log absorbance
against time was linear indicating first order in QxDC. But, the rate of the
reaction decreased with increase in the concentration of oxidant181. It is
attributed to the decrease in effective concentration of chromium (VI)
species in the reaction medium. One representation graph is given in
Fig. 4.9.
Table 4.10
[Cyclohexanol] = 2.00×10-2 mol dm-3
AcOH-H2O = 50:50 (%)
[H+] = 7.00×10-1 mol dm-3
Temperature = 313 K
[QxDC] 103
mol dm-3
2.0
k1 104
s-1
3.85
2.5
3.59
3.0
2.82
3.5
2.45
4.0
2.03
5.0
1.36
6.0
0.62
82
[QxDC] = 2.5 x 10-3 M
Fig 4.9 Plot of log absorbance versus time
83
Effect of varying the [substrate]
The rate of reaction is increased with increase the concentration of
substrate. The order with respect to substrate was found to be fractional as
evidenced by the linear plot of log k against log [s] (Fig. 4.10) with a double
reciprocal plot of k against [s] (Fig. 4.11) gave a straight line indicating
Michaelis-Menten type of kinetics in this reaction.
Table 4.11
[QxDC] = 2.50×10-3 mol dm-3
AcOH-H2O = 50:50 (%)
[H+]
Temperature = 313 K
= 7.00×10-1 mol dm-3
[Cyclohexanol] 102
mol dm-3
0.5
k1 104
s-1
1.56
1.0
2.58
2.0
3.59
3.0
4.46
4.0
5.63
5.0
6.23
6.0
7.16
84
r=0.995
B=0.59
Fig 4.10 Plot of log k versus log [substrate]
85
r=0.994
Fig 4.11 Plot of 1/k versus 1/[s]
86
Effect of varying the hydrogen ion concentration
The effect of added H+ ion on the pseudo-first order rate constant was
studied by adding perchloric acid in the region of 0.35 - 1.75 mol dm-3. The
rate of reaction increased with increase the concentration of perchloric acid.
The plot of log k against log [H+] (Fig. 4.12) gave a straight line with slope
1.44 indicating that the protonated species of the oxidant in the effective
oxidant182.
Table 4.12
[QxDC]
= 2.50×10-3mol dm-3
[Cyclohexanol] = 7.00×10-2mol dm-3
AcOH-H2O = 50:50 (%)
Temperature = 313 K
[HClO4] 101
mol dm-3
3.5
k1 104
s-1
1.13
7.0
3.59
10.5
5.99
14.0
8.58
17.5
11.78
87
Fig 4.12 Plot of log k versus log [H+]
88
Effect of varying the ionic strength
The reaction was carried out at different initial concentrations of
sodium perchlorate while the other variables were kept constant. Increase in
ionic strength of the medium by adding sodium perchlorate has no effect on
the reaction rate indicating the involvement of ion-neutral molecule in the
rate - determining step.
Table 4.13
[QxDC] = 2.50 × 10-3mol dm-3
AcOH-H2O = 50:50 (%)
[Cyclohexanol] = 7.00 × 10-2mol dm-3
Temperature = 313 K
[H+] = 7.00 × 10-1mol dm-3
[NaClO4] 102
mol dm-3
0.00
k1 104
s-1
3.59
5.05
3.52
10.10
3.61
15.15
3.63
20.20
3.54
89
Effect of varying the solvent composition
The effect of variation of solvent composition on the pseudo-first
order rate constant was also studied. The rate was found to increase when
the percent content of acetic acid increases. The plot of log k1 versus 1/D
(Fig. 4.13) gave a straight line with a positive slope. This might be probably
due to ion-dipole interaction in the rate determining step.
Table 4.14
[QxDC] = 2.50×10-3mol dm-3
[H+] = 7.00×10-1mol dm-3
[Cyclohexanol] = 2.00×10-2mol dm-3
Temperature = 313 K
AcOH-H2O
% (v/v)
40-60
D
k1 104
s-1
49.60
1.98
45-55
45.99
2.34
50-50
42.37
3.59
55-45
38.75
4.11
60-40
35.14
4.98
90
Fig 4.13 Plot of log k versus 1/D
91
Effect of added acrylonitrile
The reaction does not induce polymerization of acrylonitrile. The
added acrylonitrile has no effect on the reaction mixture indicating the
absence of free radical mechanism.
Effect of varying the [Manganous Sulphate]
The reaction was followed with varying concentration of Mn 2+ions
keeping all the other factors constant. There was an appreciable decrease in
the rate with increasing concentration of Mn2+ ions confirming the
involvement of two electron transfer process in the reaction183-185.
Table 4.15
[QxDC] = 2.50×10-3mol dm-3
AcOH-H2O = 50:50 (%)
[Cyclohexanol] = 2.00×10-2mol dm-3
Temperature = 313 K
[H+] = 7.00×10-1mol dm-3
[MnSO4] 102
mol dm-3
0.00
k1 104
s-1
3.59
0.10
2.84
0.20
2.60
0.30
2.46
0.40
2.15
92
Effect of varying the temperature
The reactions were studied in the temperature range 303 K - 333 K for
cyclohexanol keeping all the other factors constant. An increase in
temperature had resulted in an increase in the rate of the reaction. The
activation parameters have been calculated using the Eyring‟s plot186
(Fig. 4.14) and the least square analysis.
ΔH# = 39.30 k J mol-1
ΔS# = -125.06 J K-1mol-1
ΔG# = 78.44 k J mol-1 at 313 K
Table 4.16
[QxDC] = 2.50×10-3mol dm-3
[H+] = 7.00×10-1mol dm-3
[Cyclohexanol] = 2.00×10-2mol dm-3
AcOH-H2O = 50:50 (%)
Temperature
K
303
k1 104
s-1
2.56
313
3.59
323
5.04
333
6.92
93
Fig 4.14 Plot of log k2/T versus 1/T
94
4.5
Oxidation of Cyclanols by Quinoxalinium Dichromate
The kinetics of oxidation of
cyclanols viz., cyclopentanol,
cycloheptanol, cyclooctanol were carried out employing different initial
concentrations of the cyclanols. The pseudo-first order rate constant
obtained for those substrates showed fractional order dependence with
respect to cyclanols. The effect of varying substrate concentration and
temperature on rate constants is shown in Table 4.17
Table 4.17 Effect of varying the concentration of cyclanols and varying
the temperature by quinoxalinium dichromate
[QxDC]
= 2.50×10-3mol dm-3
AcOH-H2O = 50:50 (%)
[H+] = 7.00×10-1mol dm-3
Cyclanols
Cyclopentanola
Cyclopentanol b
[substrate]102
Temperature
mol dm-3
K
k1 104
s-1
313
1.0
4.24
313
2.0
5.79
313
3.0
7.34
313
4.0
8.05
313
5.0
8.91
313
6.0
9.94
303
2.0
3.47
313
2.0
5.79
323
2.0
8.37
333
2.0
11.74
Order with
respect
to substrate
0.47
-
Cyclanols
Cycloheptanola
Cycloheptanolb
Cyclooctanola
Cyclooctanolb
[substrate]102
Temperature
mol dm-3
K
k1 104
s-1
313
1.0
4.64
313
2.0
6.19
313
3.0
7.73
313
4.0
8.24
313
5.0
9.05
313
6.0
9.92
303
2.0
3.98
313
2.0
6.19
323
2.0
9.76
333
2.0
14.91
313
1.0
6.28
313
2.0
8.56
313
3.0
10.64
313
4.0
12.05
313
5.0
12.88
313
6.0
13.81
303
2.0
5.84
313
2.0
8.56
323
2.0
11.92
333
2.0
16.24
* a - Effect of varying the substrate
Order with
respect
to substrate
0.42
-
0.46
-
b - Effect of varying the temperature
96
Mechanism
From the above observations, it is clear that the reaction is showing
unit order dependence with respect to oxidant and fractional order
dependence with respect to substrate and H+ ion concentration. The reaction
is facilitated by the medium of low dielectric constant. The reaction does not
induce polymerization of acrylonitrile indicating the absence of free radical
path way. The added Mn2+ ion has a retardation on the reaction rate,
confirming the two electron transfer process involved in the reaction.
Based on the above facts the following mechanism is proposed for the
oxidation of cyclanols by quinoxalinium dichromate (Scheme 4.2).
+
K1
2(C8H8N2)Cr2O7 + H2O
O
H
O
+
O
H
OH
Cr
-
O
O
H
O
K2
-
OH
Cr
OH
Cr
O
O
+
+
2HCrO4 + C8H8N2
-
O
k3
O
(C1)
HO
+
O
OH
Cr
slow
-
O
-
O
(C1)
Scheme 4.2
O
Cr(IV)
-
O
97
Rate law:
Rate = k3C1
= k3K2 [QxDC] [s]
= k3K2K1 [QxDC] [s]
-d [QxDC]/dt = kobs [QxDC] [s]
This rate law explains all the observed experimental facts.
The thermodynamic parameters calculated for the oxidation of various
cyclanols are given in Table 4.18. The enthalpy of activation (ΔH#) values is
very low indicating a concerted mechanism as proposed. The negative
values of the entropy of activation (ΔS#) indicate extensive solvation of the
transition state over the reactants. It also reveals that the rate determining
step is less disorderly oriented relative to the reactants. Free energy of
activation (ΔG#) values is nearly constant which indicates that all the
cyclanols are oxidized by the same mechanism.
98
Table 4.18 Thermodynamic parameters for the oxidation of cyclanols by quinoxalinium dichromate
[Cyclanols] = 2.00×10-2 mol dm-3
[H+]
S.
No.
[QXDC] = 2.50 × 10-3 mol dm-3
= 7.00×10-1 mol dm-3
Cyclanols
AcOH-H2O = 50-50 (%)
k1 104
(s-1)
Order with
respect to
substrate
303 K
313 K
323 K
333 K
ΔH#
(kJ mol-1)
-ΔS#
(JK-1mol-1)
ΔG#
(kJ mol-1)
at 313 K
r
1
Cyclohexanol
0.56
2.56
3.59
5.04
6.92
10.98
197.38
72.76
0.999
2
Cyclopentanol
0.47
3.47
5.79
8.37
11.74
13.58
187.89
72.39
0.997
3
Cycloheptanol
0.42
3.98
6.19
9.76
14.91
14.94
183.08
72.23
0.999
4
Cyclooctanol
0.46
5.84
8.56
11.92
16.24
11.25
193.80
71.91
0.999
99
Applying the isokinetic relationship and using the equation163-164.
ΔH# = ΔH0# + βΔS#
ΔH# is equal to the enthalpy of activation, when ΔS#= 0 and usually
has no physical significance and β is the isokinetic temperature. A plot of
ΔH# versus ΔS# gave a straight line (Fig. 4.15). The isokinetic temperature β
obtained from the slope is 296 K. Since the β value is lower than the
experimental temperature, it indicates that this oxidation reaction is an
entropy controlled reaction187-188. This also indicates that all the cyclanols
are undergoing oxidation following a common mechanism.
100
Fig 4.15 Isokinetic plot for the oxidation of cyclanols by
quinoxalinium dichromate
(Numbers as given in Table 4.18)
101
Exner165-167 criticized the validity of such a linear correlation between
ΔH# and ΔS#, as these quantities are dependent on each other. When
measurements at two different temperatures have been made the data can be
analyzed by the following equation168-170.
log (k1)T2 = a + b log (k1)T1
where, a and b are intercept and slope and T2 > T1
The plot of log k323 K versus log k313 K (Fig. 4.16) gave a straight line
with an excellent correlation co efficient r = 0.991. Such a good correlation
indicates that all the cyclanols follow a common mechanism.
102
Fig 4.16 Exner Plot of log k323 K versus log k313 K for the oxidation of
cyclanols with quinoxalinium dichromate
(Numbers as given in Table 4.18)
103
4.6
Effect of Structure on the Reactivity of Cyclanols
The kinetic data of oxidation of the above cyclanols were analysed,
with a view to study the effect of ring size on reactivity. In general, the order
of reactivity of cyclic compounds may be of two types. In some cases, the
order of reactivity increases with the increasing size of the ring as noticed in
the oxidation of cyclic alcohols by acid bromate189 and hexacyano
ferrate (III)190. Here the order of reactivity of various cyclanols is
5 > 6 > 7 > 8, which is in conformity with Bayer‟s strain theory. During the
oxidation of these cyclanols the hydroxyl carbon atom is undergoing change
in hybridization from sp3 to sp2 to form the cyclic ketones, i.e. the
coordination number is changed from four to three. The strain which is
developed during this conversion is released as the ring size is increased.
Hence here the order of reactivity increases with the size of ring.
The reactions involving a change in coordination number from three
to four are strongly favoured in six membered ring compound relative to
5- and 7- ring derivatives, because the cyclohexanone ring with six
tetrahedral carbon atoms is highly symmetrical and stable. Moreover the
hydrogen-hydrogen repulsion are reduced to a minimum in the chair form as
a result of fully staggered constellation permitted by this form. Enlargement
of one of the ring (C-C-C) angles will decrease the symmetry, decrease the
puckering and increase the hydrogen-hydrogen repulsions. The increase in
the angle will lead to an increase in internal strain (positive strain).
Therefore in cyclohexane derivatives I-strain will oppose the reactions
involving a change in covalency of a ring atom from either four to five or
four to three. Conversely in these compounds a change in covalency of a
104
ring atom from three to four will involve a decrease in internal strain
(negative strain) and will be strongly favoured.
On the other hand 5- and 7- membered rings are appreciably strained,
primary as a result of the torsional forces about C-C single bonds. It is
proposed that the introduction of an atom with preferred 1200 angle leads to
a decrease in internal strain. Therefore, I-strain will favour the reactions
involving changes in coordination number from four to five or four to three
and oppose reactions involving covalency change from three to four.
In present study, the order of reactivity among the cyclanols are
cyclohexanol < cyclopentanol < cycloheptanol < cyclooctanol.
Small rings (3,4) are highly strained due to angle strain, whereas
normal rings (5-7), medium rings (8-11) and large rings (n>11) are said to be
relatively strain free though they have appreciable torsional strain 191.
In cyclohexane system, the reaction occurs in its flexible boat form
which has bond opposition strain though not angle strain involving four
pairs of hydrogen at the side of the boat.
In cyclopentane system, though the angle strain may not be
appreciable, there will be strain due to the eclipsing interaction with the
adjacent hydrogen atoms. The larger rate of reduction of cyclohexanone with
sodium boro hydride compared to that of cyclopentanone192 is ascribed to the
increased torsional strain caused by the conversion of sp2 to sp3
hybridisation.
105
Since in six membered ring system the conversion of sp2to sp3 is
easier because of the small bond opposition. The higher rates of
cycloheptanol and cyclooctanol are due to largest I-strain involved during
sp3to sp2 change193.
Chapter - V
SUMMARY
..
Part I
1.
The
rate
of
oxidation
of
S-phenylmercaptoacetic
acid
with
quinoxalinium dichromate under pseudo-first order conditions has been
studied at 313 K.
2.
The reaction shows first order dependence with respect to oxidant and
substrate.
3.
The reaction follows inverse first order kinetics with respect to H+.
4.
Increase in ionic strength has no effect on the reaction rate and decrease
in the dielectric constant of the medium increases the reaction rate.
5.
There is no possibility of free radical mechanism since there is no
polymerization of acrylonitrile. Increase in the concentration of
manganous sulphate retards the reaction rate which confirms the two
electrons transfer involved in the mechanism.
6.
Based on the experimental observations a suitable mechanism has been
proposed and rate law has been derived.
7.
The products of the oxidation reaction are found to be corresponding
sulfoxides.
8.
The
effect
of
various
para-
and
meta-
substituted
S-Phenylmercaptoacetic acids on the reaction rate has been studied at
four different temperatures viz., 303 K, 313 K, 323 K and 333 K and the
thermodynamic parameters are calculated using Eyring‟s equation.
107
9.
The plot of ΔH# versus ΔS# does not vary linearly, no isokinetic
relationship is observed. An excellent correlation is obtained (r = 0.999)
when log k(313 K) versus log k(303 K) values are plotted. This indicates that
all the substituents follow a common mechanism.
10. Electron releasing substituents retarded the reaction rate while the
electron withdrawing substituents enhanced the reaction rate. A plot of
log k1 versus σ was found to be linear with positive slope = +1.54.
Part II
1.
The kinetics of oxidation of cyclohexanol with quinoxalinium
dichromate in aqueous acetic acid medium has been studied at 313 K.
2.
The reaction obeys first order with respect to oxidant and fractional
order with respect to substrate. The reaction is acid catalysed by
hydrogen ions.
3.
The rate increases with decrease in the dielectric constant of the
medium and increase in ionic strength has no effect on the reaction rate.
4.
There is no appreciable change when acrylonitrile is added to the
reaction mixture indicating the absence of free radical mechanism.
5.
The rate of the reaction decreases with increase in the concentration of
manganous sulphate suggesting that the two electron transfer involved
in the reaction.
6.
A suitable mechanism has been proposed on the basis of experimental
facts and a suitable rate law is derived.
7.
The
oxidation
product
of
the
reaction
was
corresponding
cyclohexanone. The reactions are carried out at four different
108
temperatures and the thermodynamic parameters are calculated using
Eyring‟s equation.
8.
The plot of ΔH# versus ΔS# gives a straight line and from the slope the
isokinetic temperature has been calculated. However, Exner plot gives a
straight line with an excellent correlation co-efficient indicating that all
the cyclanlols follow a unified mechanism.
9.
The order of reactivity of cyclanols are Cyclohexanol < Cyclopentanol
< Cycloheptanol < Cyclooctanol.
APPENDICES
.
. .Appendix-A
Selected values of Substituent constant σ194
Substituents
Substituent constant
σm
σp
-OCH3
0.12
-0.27
-CH3
-0.07
-0.17
-Br
0.39
0.23
-Cl
0.37
0.23
-NO2
0.71
0.78
110
Appendix-B
Linear Regression Analysis
Evaluation of errors in kinetic data
In kinetic measurements as in other measurement process errors are
inevitable. A measurement process may be expected to provide meaningful
results only when it is in a statistical control throughout the measurement
process195.
In the present investigation since all the experimental data have been
analyzed by regression analysis, a brief account of the procedures employed.
Simple regression
If a series of values of a dependent variable Y (regressand) where we
suspect may be related linearly to the independent variable X (explanatory
variable or regressor), we plot Y against X by using the equation of a
straight line (I)
Y = aX +b
... (I)
Where „a‟ is the slope of the line and „b‟ is the intercept. By the
method of least squares, the value of „a‟ and „b‟ can be obtained as follows,
For „n‟ pairs of x and y values, the following equations can be defined for
convenience in calculation196.
ΣU 2 = Σ(X - X)2 = ΣX 2-n(X)2 = ΣX 2[ΣX 2-(ΣX 2)] / n
... (II)
ΣY2 = Σ( Y -Y)2 = ΣY2 - n(Y)2 = ΣY2 - [ΣY2 - (ΣY)2]/ n
…(III)
ΣUY = Σ(X- X)(Y - Y) = ΣXY-nXY = ΣXY –ΣX ΣY/ n
... (IV)
111
The slope „a‟ is given by
Slope = a = ΣUY /Σ U2
... (V)
The intercept „b‟ is given by
Intercept = b = 1/n (Σ/ - aΣX)
... (VI)
The success of the correlation is given by the standard deviation SYX
and the correlation coefficient „r‟. They are evaluated by the application of
equation (VII) and (VIII).
SYX= [ΣY2– a2ΣU2/n-2]
... (VII)
r = ΣUY/ [(ΣU2) (ΣY2)]1/2
... (VIII)
Then the variance of the regression coefficient or standard deviation
of „ a‟ is given by equation (IX).
S2aY.X = S2Y.X/ΣU2
... (IX)
SaY.X = S2Y.X/ΣU2
... (X)
Confidence limits (CL) for „a‟ are assigned in terms of „student t‟
values by equation (XI).
CL for „a‟ = a ± tαФSa.Y.X
... (XI)
where, Ф = n-2, the number of degrees of freedom and „α‟ denotes the
significant level of the „student t‟.
112
Appendix-C
Symbols and Abbreviations
[H+]
Concentration of hydrogen ion
[O]
Concentration of oxidant
[S]
Concentration of substrate
AcOH
Acetic acid
APM
2-amino 4,6-diarylpyrimidines
BAMICC
1-(benzylamino) 3-methyl imidazolium chlorochromate
BPCC
2,2′-bipyridinium chlorochromate
BTMACB
Benzyltrimethyl ammonium chlorobromate
D
Dielectric constant
Ea
Energy of activation
h
Planck‟s constant
IFC
Imidazolium fluorochromate
INDC
Isonicotinium dichromate
K
Equlibrium constant
k1
Pseudo - first order rate constant
k2
Specific reaction rate constant
kB
Boltzmann‟s constant
M
mol dm-3
MCC
Morpholinium chlorochromate
NBA
N-bromo acetamide
NBP
N-bromo phthalimide
NCN
N-chloro nicotinamide
O.D
Optical density
113
PBC
Pyridinium bromochromate
PCC
Pyridinium chlorochromate
PDC
Pyridinium dichromate
PFC
Pyridinium fluorochromate
PMA
S-Phenylmercaptoacetic acid
PtCC
Phthalazinium chlorochromate
PtDC
Phthalazinium dichromate
QBC
Quinolinium bromochromate
QFC
Quinolinium fluorochromate
QnFC
Quinaldinium fluorochromate
QxDC
Quinoxalinium dichromate
r
Correlation co-efficient
R
Gas constant
s-1
Second-1
T
Absolute temperature
TEACC
Triethyl ammonium chlorochromate
β
Isokinetic temperature
ΔG#
Gibb‟s free energy
ΔH#
Enthalpy of activation
ΔS#
Entropy of activation
λ
Wave length
ρ
Reaction constant
σ
Substituent constant
h
hour
min
minutes
mp
melting point
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152. K. G. Sekar, R. R. Muthu Chudarkodi and K. Anbarasu,
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__________
128
LIST OF PAPER PRESENTATIONS
1. 13th CRSI and 5th CRSI-RSC Symposium in Chemistry, 4-6 February
2011, NISER and KIT University, Bhuvaneswar.
K. G. Sekar and G. Manikandan
PP 62, Oxidation of S-Phenylmercaptoacetic acid by Quinoxalinium
Dichromate.
2. 14th National Symposium in Chemistry (NSC-14) and 6th CRSI-RSC
Symposium in Chemistry, 2-5 February 2012, CSIR-NIIST and
IISER, Thiruvananthapuram.
K. G. Sekar, K. Anbarasu and G. Manikandan
PP 125, Structure and Reactivity of Cyclanols towards Quinoxalinium
Dichromate Oxidation.
PP62
PP125
131
LIST OF PAPER PUBLICATIONS
1. Oxidation of substituted S-phenylmercaptoacetic
Quinoxalinium dichromate
K. G. Sekar and G. Manikandan
Oxidation Communications, 35 (3), 577 (2012).
acids
by
2. Kinetics and Mechanism of Oxidation of Cyclohexanol by
Chromium (VI)
K. G. Sekar and G. Manikandan
International Journal of Advances in Pharmacy, Biology and
Chemistry, 1(4), 450 (2012).
3. Oxidation of S-phenylmercaptoacetic acid
dichromate
K. G. Sekar and G. Manikandan
Der Chemica Sinica, 4(1), 100 (2013).
by
Quinoxalinium
4. Structure and Reactivity of Cyclanols Towards Quinoxalinium
Dichromate Oxidation
K. G. Sekar and G. Manikandan
Russian Journal of Applied Chemistry (Communicated).
Oxidation Communications 35, No 3, 577–582 (2012)
Oxidation in the presence of Cr-containing compounds
Oxidation of Substituted S-phenylmercaptoacetic Acids by
Quinoxalinium Dichromate
K. G. Sekar*, G. Manikandan
Department of Chemistry, National College, 620 001Tiruchirappalli,
Tamilnadu, India
E-mail: [email protected]; [email protected]
ABSTRACT
The conversion of S-phenylmercaptoacetic acid to the corresponding sulphoxide
was performed in 50% (v/v) water–acetic acid mixture in the presence of
perchloric acid medium. The order with respect to S-phenylmercaptoacetic acid
and quinoxa-linium dichromate were both one and inverse first order with respect
to hydrogen ion concentration. Decrease in dielectric constant of the medium
increased the rate of reaction. Ionic strength had a considerable influence on a
reaction rate, indicating the involvement of a dipole in the rate-limiting step. In
general, the electron-withdrawing substituents enhance the reaction rate and
electron-releasing substituents retard the reaction rate. A suitable mechanism and
rate law in consonance with the observed facts is proposed.
Keywords: kinetics, oxidation, phenylmercaptoacetic acid, quinoxalinium dichromate.
AIMS AND BACKGROUND
Quinoxalinium dichromate (C8H6N2H2)Cr2O7 (QxDC) has been used as a mild,
efficient and selective oxidising reagent in synthetic organic chemistry1.
H
N
2–
Cr2O7
N
H
quinoxalinium dichromate
*
For correspondence.
577
However, there are not many reports on the characteristic aspects of reactions of
QxDC studies reported so far on the kinetics of oxidation of S-phenylmercaptoacetic
2–7
acid to give diverse products, involving different intermediates in aqueous medium.
The use of an insulated acid substrate contains groups or atoms between the reaction
site and the bulk of the molecule in a similar study is rare. Now, we report the oxidation of S-phenylmercaptoacetic acid by quinoxalinium dichromate.
EXPERIMENTAL
Reagents . S-phenylmercaptoacetic acids were prepared and purified by literature
8
1
method . QxDC was prepared by a known procedure and its purity was
determined by iodometric assay. Acetic acid was refluxed over chromium trioxide
9
for 6 h and then fractionated . All other chemicals were of Analar grade. The
reaction mixture was homogeneous throughout the course of the reaction.
Kinetic measurements. The reactions were followed under pseudo-first order conditions by maintaining always the substrate concentration in excess over that of QxDC.
The reactions were carried by monitoring the decrease in the concentrations of QxDC
and were followed spectrophotometrically at 470 nm for up to 80% of the reaction. The
rate constants were evaluated from the linear plot of log absorbance against time by the
least square method and were reproducible within ± 3%.
Stoichiometry. The stoichiometric runs were carried out in the presence of excess
QxDC which reveals that 1 mol of oxidant consumes 1 mol of substrate confirming
the stoichiometry of the reaction as 1:1.
Product analysis. The kinetic reaction mixture was left to stand for 24 h under ki-netic
conditions. It was extracted with ether and the residue that separated during solvent
evaporation was analysed by IR spectroscopy. The following frequencies
–1
–1
corresponding to the sulphoxide were observed: 1024 cm (=S=O group), 1713 cm
–1
(–C=O group) and 3434 cm (–COOH group). The product was further confirmed
by TLC. The yield of sulphoxide was 90% as determined by weight measurement
of the reactant and product.
RESULTS AND DISCUSSION
The reaction was studied under different experimental conditions in the presence of
acetic acid–water (50% v/v) as solvent medium. At a constant temperature, the rate
increased steadily on increasing the concentration of the substrate as shown in Table 1.
The linear plot of lg k against lg [substrate] with a slope of unity clearly indicates that
the reaction has unit order dependence on the concentration of the substrate.
The specific reaction rate constant of k2 = k1/[S] confirms the first order in the Sphe- nylmercaptoacetic acid.
578
Table 1. Rate data on the oxidation of S-phenylmercaptoacetic acid by quinoxalinium dichromate at
313 K
3
[PMA]×10
–3
(mol dm )
2.5–12.5
5.0
5.0
5.0
5.0
3
[QxDC]×10
–3
(mol dm )
2.0
1.5–3.5
2.0
2.0
2.0
+
[H ]×10
–3
(mol dm )
3.5
3.5
3.5–17.5
3.5
3.5
AcOH:H2O
(%–v/v)
50:50
50:50
50:50
40:60–60:40
50:50
[NaClO4]×10
–3
(mol dm )
–
–
–
–
0.00–20.20
2
4
k1×10
–1
(s )
1.27–8.52
3.99–1.62
3.26–0.78
2.72–3.71
3.26–3.13
The reaction was found to be first order with respect to the oxidant as
evidenced by the good linearity in the plot of lg absorbance versus time (r=0.990).
Increase in ionic strength of the medium by adding sodium perchlorate has no
effect on the reaction rate indicating the involvement of charged species in the ratedetermining step (Table 1).
The kinetic runs were performed at different concentrations of perchloric acid
which acted as the catalyst. The rate decreased with an increase in the concentration of
+
hydrogen ion, this suggests that H ions react with S-phenylmercaptoacetic acid and
+
+
forms a non-reactive species. Plots of k versus 1/[H ] and lg k versus lg [H ] are also
straight lines with unit slope indicating an inverse first order dependence on hydrogen
ion concentration. The acetic acid composition in the solvent mixture was varied while
maintaining the other variables constant, as shown in Table 1. The rate was found to
increase considerably on increasing the acetic acid content of the medium. It is due to
the fact that the reaction is facilitated by an increase in polar-ity or nucleophilicity. The
addition of acrylonitrile, which is a very good trapper of free radicals did not have any
retarding effect on the reaction. It indicates that no free radicals participation in the
2+
reaction. The addition of Mn retard the rate of the oxidation considerably showing
that the rate-determining step involves a 2-electron transfer in the mechanism.
The reaction was performed at different temperatures, viz. 303, 313, 323 and
+
333 K while maintaining the concentrations of substrate, oxidant and H constant
(Table 2) and from the Eyring plot10 of ln (k2/T) versus 1/T, the thermodynamic parameters were calculated.
579
Table. 2. Thermodynamic parameters for the oxidation of para- and meta- S-phenylmercaptoacetic
acids by quinoxalinium dichromate
4
–1
*
S. Substitu- Order
k1 ×10 (s )
∆H
–1
(kJ
mol )
No ents
with re- 303 K 313 K 323 K 333 K
spect to
substrate
1 H
1.14
2.54 3.26 5.58 8.21 13.60
0.88
0.72 1.12 1.85 2.71 15.19
2 p-OMe
0.90
0.94 1.55 2.38 3.56 15.02
3 p-Me
1.11
5.98 9.43 13.89 19.98 13.48
4 p-Br
1.14
5.04 6.76 8.87 10.98
8.37
5 p-Cl
0.87
30.12 45.71 64.94 87.22 11.79
6 p-NO2
0.79
3.01 4.84 6.52 9.27 12.37
7 m-OMe
0.82
1.89 2.98 5.81 8.94 18.28
8 m-Me
1.08
10.21 14.45 19.21 25.09
9.73
9 m-Br
1.18
9.07 12.88 16.62 21.08
9.03
10 m-Cl
1.28
28.21 38.90 50.09 65.24
8.95
11 m-NO2
*
–∆S
–1
(J K
*
r
mol )
∆G
–1
(kJ mol )
at 313 K
184.32
183.41
182.91
181.33
196.29
180.99
187.40
169.86
191.72
194.47
190.64
71.29
72.60
72.27
70.24
69.81
68.44
71.03
71.45
69.74
69.90
68.62
0.990
0.999
0.999
0.999
0.998
0.998
0.996
0.996
0.999
0.997
0.999
–1
MECHANISM AND RATE LAW
The oxidation of S-phenylmercaptoacetic acid with quinoxalinium dichromate was
catalysed by perchloric acid. It is first order with respect to the concentrations of
+
each of the oxidant and substrate and inverse first order with respect to H . Product
analysis clearly indicates that the obtaining of the corresponding sulphoxide. From
these observations, the following mechanism and rate law were proposed.
K
1
C6H5SCH2COOH + H
(C H N )CrO H
8
7
2
2
+
+
+
C6H5SC H3COOH
K2
(C H N)Cr O + H
7
8
7
2
2
+
7
K3
+
(C8H7N2)Cr2O7H + C6H5SCH2COOH
complex
O
k
||
4
+
complex C H –S– –––––→ CH COOH + H + Cr(IV)
6
5
2
Rate law:
rate = k4 [complex]
+
= k4 K3 [QxDCH ] [PMA]
+
+
= k4 K3 K2 [QxDC] [PMA] [H ]/K2 [H ]
+
+
–d(QxDC)/dt = k4K1 K2 K3 [QxDC] [PMA] [H ]/K2 [H ]
The proposed mechanism and the rate law support all the observations made
including the effect of solvent polarity and the negative entropy of activation.
580
EFFECT OF SUBSTITUENTS
The rate constant k1 was estimated for the substituted S-phenylmercaptoacetic acids at
4 different temperatures, viz. 303, 313, 323 and 333 K. The thermodynamic parameters
have been computed from a plot of ln k2/T versus 1/T using the Eyring equation. The
*
negative values of the entropies of activation (∆S ) suggested that the transition state
formed was considerable rigid, resulting in a reduction in the degree of freedom of the
*
molecules. The constancy of the (∆G ) values indicated a common mechanism for the
*
*
oxidation of all the substrates. As ∆H and ∆S do not vary linearly, no isokinetic
relationship is observed. This indicated the absence of enthalpy – entropy compensa11
12
*
tion effect . Exner criticised the validity of such a linear correlation between ∆H
*
and ∆S as the quantities are dependent on each other, when measurements at 2 different temperatures have been made. The rate data can be analysed by the following
13
equation :
lg k1 (T2) = a + b lg k1 (T1),
where a and b are intercept and slope and T2>T1.
The plot of lg k1 (303 K) against lg k1 (313 K) gave a straight line with r =
0.997 – such a good correlation indicates that the oxidation of the substrates with
different substituents follows a common mechanism.
To have an idea about the order with respect to each of the substrate the
oxidation has been studied at 313 K and the results are given in Table 2. It is
interesting to note that all the substituted S-phenylmercaptoacetic acids show a unit
order dependence on the reaction rate.
The rate data for the oxidation of para- and meta-substituted S-phenylmercaptoacetic acids give a good correlation for the plot of lg k versus σ (Fig. 1) (r = 0.994, ρ
= +1.54) with the Hammett value at 313 K. Similar phenomenon has been observed in
the oxidation of substituted benzaldehydes by PFC (Ref. 14) and IDC (Ref. 15). The
positive ‘ρ’ value indicates that electron-withdrawing substituents enhance the rate of
oxidation and electron-releasing substituents decrease the rate of the reaction.
Fig. 1. The Hammett plot of lg k313 K versus σ for the oxidation of S-phenylmercaptoacetic acids by
QxDC (numbers as given in Table 2)
581
CONCLUSIONS
The oxidation of S-phenylmercaptoacetic acid by quinoxalinium dichromate was
+
studied in full depth and a mechanism involving the substrate, oxidant and H is
proposed. In the slow rate-determining step, the substrate reacts with the positively
charged species. The product is the corresponding sulphoxide. The orders with respect
to the concentrations of substrate and oxidant are one. The perchloric acid reacts with
substrate to form a non-reactive species. The negative sign of the entropy change suggests that the transition state is more orderly when compared with the reactants.
REFERENCES
1. N. DEGRIMENBASI, B. OZGUN: Quinoxalinium Dichromate: A New and Efficient Reagent for
the Oxidation of Organic Substrate. Monatshefte fur Chemie, 133, 1417 (2002).
2. S. KABILAN, M. UMA, K. KRISHNASAMY, P. SANKAR: Oxidation of S-phenylmercaptoacetic
Acid by Pyridinium Dichromate. J. Indian Council Chem., 10 (1), 21 (1994).
3. S. KABILAN, K. KRISHNASAMY, P. SANKAR: Oxidative Cleavage of Phenylthioacetic Acid
by Pyridinium Dichromate in Acetonitrile Medium: Kinetic and Correlation Study. Oxid
Commun, 18 (3), 288 (1995).
4. S. KABILAN, R. GIRIJA, V. RAJAGOPAL: Oxidative Cleavage of S-phenylmercaptoacetic
Acid by Pyridinium Chlorochromate: Kinetic and Correlation Analysis. Int. J. Chem. Kinet., 31
(10), 109 (1999).
5. K. SATHIYANARAYANAN, R. SUSEELA, CHANG WOO LEE: Oxidation of S-phenylmercaptoacetic Acid by N-chloronicotinimide: A Kinetic Study. J. Ind. Eng. Chem., 12 (2), 280 (2006).
6. K. SATHIYANARAYANAN, C. PAVITHRA, CHANG WOO LEE: Kinetics and Mechanism of
S-phenylmercaptoacetic Acid by Chromium(VI). J. Ind. Eng. Chem., 12 (5), 727 (2006).
7. N. M. I. ALHAJI, A. M. UDUMAN MOHIDEEN, K. KALAIMATHI: Mechanism of Oxidation
of p-substituted Phenylthio Acetic Acids with N-bromophthalimide, E. J. Chem., 8 (1), 1 (2011).
8. S. GABRIEL: Oxidation of Some Sulphur Compounds. Ber., 12, 1879 (1939).
9. K. S. P. ORTON, A. E. BRADFIELD: The Purification of Acetic Acid. The Estimation of Acetic
Anhydride in Acetic Acid. J. Chem. Soc., 983 (1927).
10. H. EYRING: The Activated Complex in Chemical Reactions. J. Chem. Phys., 3, 107 (1935).
11. M. G. ALDER, J. E. LEFFLER: The Role of the Solvent in Radical Composition Reactions:
Phenyl Azotriphenyl Methane. J. Am. Chem. Soc., 76, 1425 (1954).
12. O. EXNER: Concerning the Isokinetic Relationship. Nature, 201, 488 (1964).
13. M. J. MALAWSKI: The Linear Relation between Enthalpy and Entropy of Activation. Roczniki
Chem., 38, 1129 (1964).
14. P. S. RAMAKRISHNAN, P. CHOCKALINGAM: Kinetics of Oxidation of Substituted
Benzalde-hydes by Pyridinium Fluorochromate in Acetic Acid – Perchloric Acid Medium. J.
Indian Chem. Soc., 70, 581 (1993).
15. K. BALASUBRAMANIAN, K. LAKSHMANAN, K. G. SEKAR: Kinetics and Mechanism of Oxidation of Aromatic Aldehydes by Imidazolium Dichromate. Asian J. Chem., 11 (4), 1451 (1999).
Received 14 April 2012
Revised 17 May 2012
582
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IJAPBC – Vol. 1(4), Oct- Dec, 2012 ISSN: 2277 - 4688
__________________________________________________________________________________
INTERNATIONAL JOURNAL OF ADVANCES IN PHARMACY,
BIOLOGY AND CHEMISTRY
Research Article
Kinetics and Mechanism of Oxidation of
Cyclohexanol by Cr (VI)
KG. Sekar* and G. Manikandan
Department of Chemistry, National College, Tiruchirappalli, Tamil Nadu, India.
ABSTRACT
The kinetics of oxidation of a cyclohexanolwith quinoxalinium dichromate (QxDC) has been investigatedin
aqueous acetic acid 50% (v/v) solution in the presence of perchloricacid. The reaction is first order with
respect to oxidant QxDC and exhibits Michaelis- Menton dependence on substrate concentration. The rate of
+
reaction increased with increase the concentration of perchloric acid and the order with respect to [H ] was
found to be fractional. From the kinetic data obtained, the activation parameters have been calculated
and a plausible mechanism has been proposed.
Keywords: Oxidation, Cyclohexanol, Quinoxalinium dichromate, Kinetics.
INTRODUCTION
Quinoxalinium dichromate (QxDC), one of the Cr (VI) compounds in reported to be a neutral and mild oxidant
for selective oxidation.
H
N
2-
Cr2O7
N
H
Quinoxalinium dichromate
1-2
MATERIALS AND
METHODS Materials
The cyclohexanolare purchased from Aldrich
chemicals, QxDCwas prepared by the literature
1
method and its purity was checked by estimating
Cr(VI) iodometrically. Acetic acid (AnalaR) was
refluxed over CrO3and distilled. All other
chemicals used were AnalaR grade. The reaction
mixture was homogeneous throughout the course
of the reaction.
Kinetics of oxidation of some organic substrates
by Quinoxalinium dichromate has already been
reported. A survey into the literature on the
kinetics of oxidation of cyclanols with various
oxidant shows that the reactivity varies with the
3-12
type of oxidant
. The difference in the reactivity
has been explained by the I-strain theory. The
present work on the oxidation of cyclohexanolby
QxDC is to ascertain the nature and the order of
reactivity of these compounds under the given
kinetic conditions.
450
IJAPBC – Vol. 1(4), Oct- Dec, 2012 ISSN: 2277 - 4688
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__________________________________________________________________________________
Kinetic Measurements
All the reactions were carried out in blackcoated
vessels to avoid the possible photochemical
reactions if any. The kinetic measurements were
carried out using spectrophotometer (Systronics) at
470nm.All kinetic runs were made in aqueous
acetic acid 50% (v/v) under pseudo-first order
conditions by keeping the substrate always in
excessover that of oxidant. The rate constants were
evaluated from the linear plot of log absorbance
against time by the least square method. The
results were reproducible within ±3%error.
Effect of substrate
The rate of reaction is increased with increase the
concentration of substrate. The order with respect
to substrate was found to be fractional as
evidenced by the linear plot of log k against log [s]
with a double reciprocal plot of k against [s] gave
a straight line indicating Michaelis-Menton type of
kinetics in this reaction.
+
Effect of H ion
+
The effect of added H ion on the pseudo-first
order rate constant was studied by adding HClO4
-3
in the region of 0.35-1.75 mol dm .The rate of
reaction increased with increase the concentration
+
of HClO4.The plot of log k against log[H ] give a
straight line with slope 0.503(Fig. 1) indicating
that the protonated species of the oxidant in the
effective oxidant. It can be concluded that the
15
reaction is simply an acid catalyzed one .
Product Analysis
The same experimental conditions were used for the
kinetic determinations;solution of oxidant (0.10 mol),
cyclohexanol (0.12 mol) and perchloric acid were
mixed and kept under nitrogen atmosphere for 24h,
for the completion of the reaction. The products were
extracted with chloroform and the organic layer
washed with water dried over anhydrous sodium
sulphate and then concentrated, the products are
separated by column chromatography using silica gel
and eluting with varying proportions (100:0 to 70:30)
(v/v) of hexane and chloroform, and were identified
after concentrating the different fractions, the
products
Effect of solvent andIonic strength
The effect of variation of solvent composition on
the pseudo-first order rate constant was also
studied. The rate was found to increase when the
percent content of acetic acid increases. The data
in Table 1 shows that the influence of ionic
strength on rate constant is not significant.
The reaction mixture showing the absence of any
free radical in the reaction has ruled out the
possibility of a one electron transfer during the
addition of acrylonitrile. But a noticeable catalytic
effect on the reaction rate on the addition of
MnSO4.
were also detected by IR spectral studies and spot
13
tests .
RESULTS AND DISCUSSION
Oxidation of cyclohexanol
The detailed kinetic data on the oxidation of
cyclohexanol is given in Table1.
Effect of oxidant
+
At fixed [H ] with [substrate] in excess, the plot of
log absorbance againsttime was linear indicating
first order in QxDC. But, the rate of the reaction
decreased with increase in the concentration of
14
oxidant .It is attributed to the decrease in
effective concentration of Cr (VI) species in the
reaction medium.
Effect of Temperature
The reactions were studied in the temperature range
303 K – 333 K for cyclohexanol (Table 2). An
increase in temperature had resulted in an increase in
the rate of the reaction. The activation parameters
16
have been calculated using the Eyring’s plot and the
least square analysis. From the kinetic data the
following mechanism has been proposed.
Table 1: Rate data on the oxidation of cyclohexanol by Quinoxalinium dichromate at 313 K
[Cyclohexanol] 10
-3
(mol dm )
1.0 - 5.0
2.0
2.0
2.0
2.0
2
[QxDC] 10
3
-3
(mol dm )
2.5
2.0 - 4.0
2.5
2.5
2.5
+
2
[H ] 10
AcOH:H2O
[NaClO4] 10
(mol dm )
7.0
7.0
3.5 - 17.5
7.0
7.0
(v/v)
50 : 50
50 : 50
50 : 50
40:60 - 60:40
50 : 50
(mol dm )
0.00 - 20.20
-3
-3
k1 x 10
4
-1
(s )
2.58 - 6.23
3.85 - 2.03
2.53-5.82
1.98-4.98
3.59-3.54
Table 2: Effect of temperature on the oxidation of cyclohexanol by quinoxalinium dichromate
-2
-3
-3
-3
[Cyclohexanol] = 2.0 x 10 (mol dm )
[QxDC] = 2.5 x 10 (mol dm )
+
-3
[H ] = 0.7 (mol dm )
AcOH:H2O(v/v) = 50:50
4 -1
S.No.
Temperature K
k1 x 10 (s )
1
303
2.85
2
313
3.59
3
323
4.56
4
333
5.28
#
-1
#
-1
-1
#
-1
H = 10.98 kJmol S = 197.38 JK mol G = 72.76 kJmol
451
IJAPBC – Vol. 1(4), Oct- Dec, 2012 ISSN: 2277 - 4688
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__________________________________________________________________________________
Mechanism and Rate Law
(C8H8N2)Cr2O7
O H
2-
K1
+ H2O
O
-
2HCrO4 + C8H8N2
OH
O H
O
OH
K2
+
Cr
-
O H
Cr
O
O
O
OH
-
O
O
O
HO
OH
k3
Cr
+
Cr
slow
O
O
-
-
O
OCr (IV)
Rate = k3C1
= k3K2 [QxDC] [S]
= k3K2K1 [QxDC] [S]
-d [QxDC]/dt = kobs[QxDC][S]
r = 0.997
B = 0.503
+
Fig. 1: Plot of log k against log [H ]
2. Ozgun B and Degirmenbasi N. Oxidation
of substituted Benzyl Alcohols by
Quinoxalinium dichromate – A Kinetic
Study, Monat. Fur Chemie. 2004;135:483491.
3. Corey
EJ
and
Suggs
JW.
PyridiniumChlorochromate: An Efficient
REFERENCES
1. Degirmen basi N and Ozgun B.
Quinoxalinium Dichromate: A New and
Efficient Reagent for the Oxidation of
Organic Substrate,Monat. Fur Chemie.
2002;133:1417-1421.
452
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__________________________________________________________________________________
Reagent for Oxidation of Primary and
Secondary
Alcohols
to
Carbonyl
Compounds,
Tetrahedron
Lett.
1975;16(31):2647-2650.
4. Corey EJ and Schmidt TG. Asymmetric
Introductions in Phase Transfer Catalyzed
Reactions: A Comment on a Structural
Feature of the Catalyst, Tetrahedron Lett.
1979;20(5):403-404.
5.
6.
9. Gurumurthy R, Karthikeyan B and
Selvaraju M. Reactivity of cyclanols
towards quinolinium chlorochromate
oxidation.
Oxid
Commun.
1999;22(1):103-106.
10. Sekar KG, Muthuchudarkodi RR and
Anbarasu K. Oxidation of Cyclanols with
Quinolinum Dichromate – A Kinetic
Study. Oxid Commun. 2007;30(2):391397.
11. Sekar KG and Anbarasu K.Structure and
Reactivity
of
Cyclanols
towards
Nicotinium Dichromate Oxidation. Oxid
Commun. 2008;31(1):199-203.
12. Sekar KG and Prabakaran A. Structure
and Reactivity of Cyclic Alcohols
Towards Pyrazinium Chlorochromate
Oxidation. Oxid Commun. 2008;31(2):
348-355.
13. Feigl F. Spot Test in Organic Analysis.
Elsevier, Amsterdam. 1966;482.
14. Krishnapillay M and Thirunavukkarasu A.
Kinetics of oxidation of some substituted
piperidones by acid permanganate. Indian
J Chem. 1981;20B:583 – 585
15. Ravishankar M, Sekar KG and
Palaniappan AN. Kinetic studies on the
oxidation of somepara- and metaSubstituted
phenols
byquinolinium
dichromate,Afinidad, 1998;477:357-362.
16. Eyring H. The activated complex in
chemical reactions. J Chem Phy.
1935;33:107- 114.
Bhattacharjee MN, Chaudri MK and
Purkayastha S. Some Aspects of Pyridinium
Fluorochromate Oxidations – Stoichiometry
of Oxidation of Alcohols, Evidence for
Oxygen Transfer and the Identity of the
Reduced Chromium Species, Tetrahedron.
1987;43:5389-5392.
Ganapathy K, Gurumurthy R, Mohan N and
Sivagnam G. Kinetics and Mechanism of
Oxidation of Cyclohexanol by 1Chlorobenzotriazole in Acid Medium,
Monat. Fur Chemie. 1987;118:583-587.
7. Agarwal S, Tiwari HP and Sharma JP.
Pyridinium
Chlorochromate:
An
Improved Method for its Synthesis and
Use of Anhydrous Acetic Acid as Catalyst
for Oxidation Reactions, Tetrahedron.
1990;46:4417-4420.
8. Chaudri MK, Chettri SK, Lyndem S, Paul
PC and Srinivas P. Quinolinium
Fluorochromate (QFC): An Improved Cr
(VI) Oxidant for Organic Substrates, Bull.
Chem Soc Jpn. 1994;67:1894-1898.
453
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Pelagia Research Library
Der Chemica Sinica, 2013, 4(1):100-104
ISSN: 0976-8505
CODEN (USA) CSHIA5
Oxidation of S-phenylmercaptoacetic acid by quinoxalinium dichromate
K. G. Sekar* and G. Manikandan
Department of Chemistry, National College, Tiruchirappalli 620 001, Tamilnadu, India.
_____________________________________________________________________________________________
ABSTRACT
The conversion of S-phenylmercaptoacetic acid to the corresponding sulfoxide was performed in 50% (v/v) wateraceticacid mixture in the presence of perchloric acid medium. The order with respect to S-phenylmercaptoacetic acid
and quinoxalinium dichromate were both one and inverse first order with respect to hydrogen ion concentration.
Decrease in dielectric constant of the medium increased the rate of reaction. Ionic strength had a considerable
influence on a reaction rate, indicating the involvement of a dipole in the rate-limiting step. A suitable mechanism and
rate law in consonance with the observed facts is proposed.
Key words: kinetics, oxidation, s-phenylmercaptoacetic acid, quinoxalinium dichromate
_____________________________________________________________________________________________
INTRODUCTION
Quinoxalinium Dichromate (C8H6N2H2)Cr2O7(QxDC) has been used as a mild, efficient and selective oxidising
reagent in synthetic organic chemistry [1].
Quinoxalinium dichromate
However, there are not many reports on the characteristic aspects of reactions of QxDC studies reported so far on the
kinetics of oxidation of S-phenylmercaptoacetic acid [2-7] to give diverse products, involving different intermediates in
aqueous medium. The use of an insulated acid substrate contains groups or atoms between the reaction site and the bulk
of the molecule in a similar study is rare. Now, we report the oxidation of S-phenylmercaptoacetic acid by
quinoxalinium dichromate.
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MATERIALS AND METHODS
Reagent: S-phenylmercaptoacetic acid was prepared and purified by literature method [8]. QxDC was prepared by a
known procedure [1] and its purity was determined by iodometric assay. Acetic acid was refluxed over chromium
trioxide for 6 h and then fractionated [9]. All other chemicals were of AnalaR grade. The reaction mixture was
homogeneous throughout the course of the reaction.
Kinetic measurements: The reactions were followed under pseudo-first order conditions by maintaining always the
substrate concentration in excess over that of QxDC. The reactions were carried by monitoring the decrease in the
concentrations of QxDC and were followed spectrophotometrically at 470 nm for up to 80% of the reaction. The rate
constants were evaluated from the linear plot of log absorbance against time by the least square method and were
reproducible within ± 3%.
Stoichiometry: The stoichiometric runs were carried out in the presence of excess QxDC which reveals that one mole
of oxidant consume one mole of substrate confirming the stoichiometry of the reaction as 1:1.
Product analysis: The kinetic reaction mixture was left to stand for 24 h under kinetic conditions. It was extracted with
ether and the residue that separated during solvent evaporation was analyzed by IR Spectroscopy. The following
-1
-1
frequencies corresponding to the sulfoxide were observed: 1024 cm (=S=O group), 1713 cm (-C=O group) and
-1
3434 cm (-COOH group).The product was further confirmed by TLC. The yield of sulfoxide was 90% as determined
by weight measurement of the reactant and product.
RESULTS AND DISCUSSION
The reaction was studied under different experimental conditions in the presence of acetic acid–water (50% v/v) as
solvent medium. At a constant temperature, the rate increased steadily upon increasing the concentration of the
substrate as shown in Table 1. A linear plot of log k against log [substrate] with a slope of unity (Fig 1). It is clear that
indication of the fact that the reaction has unit order dependence on the concentration of the substrate. The specific
reaction rate constant of k2 = k1/[s] confirms the first order in the S-phenylmercaptoacetic acid.
The reaction was found to be first order with respect to the oxidant as evidenced by a good linearity in the plot of log
absorbance versus time (r = 0.990). Increase in ionic strength of the medium by adding sodium perchlorate has no
effect on the reaction rate [10-11] indicating the involvement of charged species in the rate determining step (Table 1).
The kinetic runs were performed at different concentrations of perchloric acid which acted as the catalyst. The rate
+
decreased with an increase in the concentration of hydrogen ion, this suggests that H ions react with S+
+
phenylmercaptoacetic acid and form a non-reactive species. Plot of k versus 1/[H ] and log k versus log [H ] are also
straight line with unit slope indicating an inverse first order dependence on hydrogen ion concentration. The acetic acid
composition in the solvent mixture was varied while maintaining the other variables constant, as shown in Table 2. The
rate was found to increase considerably upon increasing the acetic acid content of the medium. It is due to the fact that
the reaction is facilitated by an increase in polarity or nucleophilicity. The addition of acrylonitrile, which is a very
good trapper of free radicals, did not have any retarding effect on the reaction. It indicates that no free radicals
2+
participation in the reaction [12-13]. The addition of Mn retard the rate of the oxidation considerably showing that
the rate determining step involves a two-electron transfer in the mechanism [14-16].
The rate data in Table 3 showed that the reaction was performed at different temperatures viz., 303 K, 313 K, 323 K
+
and 333 K while maintaining the concentrations of substrate, oxidant and H constant. From the Eyring’s plot [17] of ln
(k2/T) versus 1/T, the thermodynamic parameters were calculated.
MECHANISM AND RATE LAW
The oxidation of S-phenylmercaptoacetic acid with quinoxalinium dichromate catalysed by perchloric acid. It is first
+
order with respect to the concentrations of each of the oxidant and substrate and inverse first order with respect to H .
Product analysis clearly indicates that the corresponding sulfoxide. From these observations, the following mechanism
and rate law were proposed.
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The proposed mechanism and the rate law support all the observations made including the effect of solvent polarity and
the negative entropy of activation.
Table.1 Rate data on the oxidation of S-phenylmercaptoacetic acid by quinoxalinium dichromate at 313 K in 50% aqueous acetic
acid medium
2
[PMA] 10
-3
(mol dm )
0.25
0.50
0.75
1.00
1.25
0.50
0.50
0.50
0.50
0.50
0.50
0.50
0.50
0.50
0.50
[QxDC] 10
-3
(mol dm )
0.20
0.20
0.20
0.20
0.20
0.15
0.20
0.25
0.30
0.35
0.20
0.20
0.20
0.20
0.20
2
+
[H ]
-3
(mol dm )
0.35
0.35
0.35
0.35
0.35
0.35
0.35
0.35
0.35
0.35
0.35
0.70
1.05
1.40
1.75
4
k1 10
-1
(s )
1.27
3.26
4.42
6.24
8.52
3.99
3.26
2.35
1.97
1.62
3.26
2.55
2.03
1.36
0.78
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Table. 2 Rate data on the oxidation of S-phenylmercaptoacetic acid by quinoxalinium dic hromate at 313 K
2
2
4
AcOH: H2O [NaClO4] 10
[MnSO4] 10
k1 10
-3
-3
-1
(%-v/v)
(mol dm )
(mol dm )
(s )
40 : 60
2.72
45 : 55
2.99
50 : 50
3.26
55 : 45
3.48
60 : 40
3.71
50 : 50
5.05
3.30
50 : 50
10.10
3.18
50 : 50
15.15
3.22
50 : 50
20.20
3.13
50 : 50
0.15
1.78
50 : 50
0.30
1.72
50 : 50
0.45
1.74
50 : 50
0.60
1.69
-2
-3
-2
-3
+
-3
[PMA] = 0.5 0 x 10 mol dm [QxDC] = 0.20 x 10 mol dm [H ] = 0.35 mol dm
Table 3.Thermodynamic param eters for the oxidation of s-phenylmercaptoacetic acid by quin oxalinium dichromate
Temperature
(K)
303
313
323
333
-2
k1 10
-1
(s )
2.54
3.26
5.58
8.21
4
Thermodynamic and
Activation Parameters
#
∆H = 13.60 kJmol
#
-3
-1
-1
-1
∆S = -184.32 JK mol
#
-1
∆G = 71.29 kJmol
-1
Ea = 16.20 kJmol
-2
-3
[ PMA] = 0.50 x 10 mol dm [QxDC] = 0.20 x 10 mol dm
+
-3
[H ] = 0.35mol dm AcOH : Water = 50:50 (%- v/v)
Fig.1. Plot of log k1versus log [s]
CONCLUSION
Oxidation of S-phenylmercaptoacetic acid by quinoxalinium dichromate was studied in full depth and a mechanism
+
involving the substrate, oxidant and H is proposed. In the slow rate-determining step, t he substrate reacts with the
positively charged species. The pro duct is the corresponding sulfoxide. The orders with r espect to the concentrations
of substrate and oxidant are one. The perchloric acid reacts with substrate to form a non-reactive species. The negative
sign of the entropy chan ge suggests that the transition state is more orderly when compared with the reactants.
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[3] S. Kabilan, K. Krishnasamy, P. Sankar, Oxid. Commun., 1995, 18(3), 288.
[4] S .Kabilan, R. Girija, V .Rajagopal, Int. J. Chem. Kinet., 1999, 31(10), 109.
[5] K. Sathiyanarayanan, R. Suseela, Chang Woo Lee, J. Ind. Eng. Chem., 2006, 12(2), 280.
[6] K. Sathiyanarayanan, C. Pavithra, Chang Woo Lee, J. Ind. Eng. Chem., 2006, 12(5), 727.
[7] N. M. I. Alhaji, A. M .UdumanMohideen, K. Kalaimathi, E. J. Chem., 2011, 8(1), 1.
[8] S.Gabriel, Ber.,1939, 12, 1879.
[9] K. S .P. Orton, A. E. Bradfield, J. Chem. Soc., 1927, 983.
[10] R. A. Singh, Kaminisingh, S. K. Singh, J.Chem. Pharm. Res.,2010,2(3),684
[11] K. G. Sekar, R. V. Sakthivel, J. Chem. Pharm. Res.,2012, 4(7),3391
[12] J. S. Littler, W. A. Waters, J. Chem. Soc.,1959,1299.
[13] K. G. Sekar, M. Vellaisamy, Der Chemica Sinica., 2012,3(3),703
[14] S. Banfi, M. Cavazzini, G.Pozzi,S.V. Barkanova, O.L.Kaliya., J.Chem.Soc.,Perkin Trans.,2000,2,879.
[15] K. G. Sekar, S.K. Periyasamy,J. Chem. Pharm.Res.,2012,4(4),2153.
[16] Firoz Ahmad, Ritu Singh, M.Abbos Siddiqui,J. Chem.Pharm.Res.,2012,4(1),608.
[17] H. Eyring, J. Chem. Phys., 1935, 3, 107.
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