Alkaline hexacyanoferrate(III) oxidation of substituted 4-oxo acids: A mechanistic study

Article abstract of DOI:10.1002/poc.968

The kinetics of the oxidation of substituted 4-oxo-4-arylbutanoic acids by hexacyanoferrate(III) in aqueous alkaline medium were studied. The reactions are found to be first order in each of the reagents, the oxo acid, hydroxide ion and the hexacyanoferrate(III) ion. A mechanism involving the formation of enolate anion from the oxo compound and subsequent rate-determining electron transfer is proposed. Application of the Olson-Simonson rule substantiated the participation of negatively charged ions in the rate-determining step. The experimental rate constant is found to be composite, being the product of the equilibrium constant for the enolate anion formation and the rate constant of the oxidation step (Kk₂). The effect of addition of solvents such as methanol and DMSO is discussed and a quantitative relationship is derived between the ionic strength and the composite rate constant. The low activation energies and even the negative activation energy in the oxidation of 4-(3-nitropheyl)-4-oxobutanoic acid are explained. The ρ values are positive and decrease with temperature. Contrary to the expected electronic effect, the methyl group in the ortho position accelerates the reaction owing to steric factors. From the intersection of the lines in the Hammett and Arrhenius plots, the isokinetic relationship is discussed. Copyright

Full text of DOI:10.1002/poc.968

JOURNAL OF PHYSICAL ORGANIC CHEMISTRY
J. Phys. Org. Chem. 2005; 18: 1042–1049
Published online 7 July 2005 in Wiley InterScience (www.interscience.wiley.com). DOI: 10.1002/poc.968
Alkaline hexacyanoferrate(III) oxidation of substituted
4-oxo acids: a mechanistic study
F. J. Maria Pushparaj,1 S. Kannan,2 L. Vikram,1 Lalitha S. Kumar2 and K. S. Rangappa3*
1Department of Chemistry, Government Arts College, Udhagamandalam 643002, India
2Department of Chemistry, Indira Gandhi National Open University, New Delhi, India
3Department of Studies in Chemistry, University of Mysore, Mysore 570006, India
Received 5 March 2005; revised 6 May 2005; accepted 12 May 2005
ABSTRACT: The kinetics of the oxidation of substituted 4-oxo-4-arylbutanoic acids by hexacyanoferrate(III) in
aqueous alkaline medium were studied. The reactions are found to be first order in each of the reagents, the oxo acid,
hydroxide ion and the hexacyanoferrate(III) ion. A mechanism involving the formation of enolate anion from the oxo
compound and subsequent rate-determining electron transfer is proposed. Application of the Olson–Simonson rule
substantiated the participation of negatively charged ions in the rate-determining step. The experimental rate constant
is found to be composite, being the product of the equilibrium constant for the enolate anion formation and the rate
constant of the oxidation step (Kk2). The effect of addition of solvents such as methanol and DMSO is discussed and a
quantitative relationship is derived between the ionic strength and the composite rate constant. The low activation
energies and even the negative activation energy in the oxidation of 4-(3-nitrophey1)-4-oxobutanoic acid are
explained. The ꢀ values are positive and decrease with temperature. Contrary to the expected electronic effect, the
methyl group in the ortho position accelerates the reaction owing to steric factors. From the intersection of the lines in
the Hammett and Arrhenius plots, the isokinetic relationship is discussed. Copyright # 2005 John Wiley & Sons, Ltd.
KEYWORDS: 4-oxo acids; oxidation; hexacyanoferrate(III); mechanism; kinetics
INTRODUCTION
picture of the reaction and also to seek an explanation for
the above-mentioned features of oxidation by hexacya-
noferrate. The solubility in an aqueous alkaline medium
due to the carboxyl group makes these oxo acids a system
of choice.
Hexacyanoferrate(III) is a complex ion, capable of ab-
stracting one electron from an electron-rich site in an
organic molecule. The transfer of the electron occurs
from the substrate to the metal ion by an outer-sphere
mechanism.1 It is capable of oxidizing under both acidic
and alkaline conditions and has been utilized extensively
in the oxidation of organic compunds.2 The oxidation of
organic compounds by hexacyanoferrate(III) in alkaline
medium presents some special kinetic features. The rate
constant decreases when the initial concentration of the
oxidant3–6 increases. Although the ionic strength of the
medium has a great influence on the rate of oxidation,7–9
no quantitative relationship has so far been established
between them. The experimental energy of activation is
low in all cases and no explanation has so far been
offered. The Hammett reaction constant ꢀ for the oxida-
tion of aromatic oxo compounds such as acetophenone
has high positive values,10 although oxidation reactions,
in general, are expected to yield low negative values.
The present study on the oxidation of 4-oxo-4-arylbu-
tanoic acids was undertaken to gain a clear mechanistic
EXPERIMENTAL
Preparation of 4-oxo-4-acids
4-Oxo-4-arylbutanoic acids were prepared by standard
procedures found in the literature.11–14 These involve
mainly Friedel–Crafts reaction of benzene and its
analogues with succinic anhydride in the presence of
anhydrous aluminium chloride. Nitration of 4-oxo-
4-arylbutanoic acids gave the 3-nitro compounds.15 4-
Oxo-2,4-diphenylbutanoic was prepared from chalcone
via the formation of the nitrile16 and subsequent hydro-
lysis. All the oxo acids were crystallized twice from
water and their purity was checked by their melting-
points and UV, IR and NMR spectra.
Kinetic measurements
*Correspondence to: K. S. Rangappa, Department of Studies in
Chemistry, University of Mysore, Mysore 570006, India.
E-mail: rangappaks@chemistry.uni-mysore.ac.in
All the kinetic measurements were carried out using
doubly distilled water with analytical-grade KCl,
Contract/grant sponsor: UGC SAP (Phase I)-DRS.
Copyright # 2005 John Wiley & Sons, Ltd.
J. Phys. Org. Chem. 2005; 18: 1042–1049

ALKALINE HEXACYANOFERRATE(III) OXIDATION OF 4-OXO ACIDS
1043
K3Fe(CN)6 and NaOH. In a typical run, 20 cm3 of freshly
prepared 0.05 M oxo acid in 0.1 M NaOH was pipetted into
a reaction bottle that was painted black to avoid photo-
chemical conversion of hexacyanoferrate(II) to hexacya-
noferrate(III). A 25 cm3 volume of water was added and
the bottle was thermostated. The reaction was started by
pipetting 5 cm3 of thermostated 0.02 M hexacyanoferra-
te(III) solution. The reaction was followed by estimating
the unreacted hexacyanoferrate(III) iodimetrically.
Addition of 5 cm3 at 10% ZnSO4 solution ensured the
completion of the reaction between iodide and hexacya-
noferrate(III) ions. The pesudo-first-order rate constants
were calculated by a graphical method and the values
were reproducible to within 3%. The ionic strength of the
medium was maintained using KCl.
clear picture, the oxidation of 4-oxo-2,4-diphenylbuta-
noic acid was carried out. It was found that as soon as the
oxo acid solution was mixed with the oxidant solution, a
precipitate was formed that was identified as chalcone by
its melting-point and IR spectrum.
The stoichiometry of the reaction (substrate:oxidant)
was observed to be variable and ranged from 1:2 to 1:5.
The ultimate products of oxidation of the oxo acid are
benzoic acid and CO2 and this oxidation requires 6 equiv.
of the oxidant, whereas oxidative decarboxylation would
require 2 equiv. The observed stoichiometry could there-
fore arise from the interplay of several such incomplete
oxidations.
Duplicate experiments were conducted in an atmo-
sphere of nitrogen and without nitrogen. Since no differ-
ence was observed in the results, all the experiments
reported in this paper were obtained without nitrogen.
Freshly prepared solutions of oxo acids in sodium hydro-
xide were used to avoid any possible side-reactions.
Polymerization test
When a few drops of methyl acrylate were added to the
oxo acid solution in NaOH containing hexacyanoferra-
te(III), turbidity occurred, indicating polymerization.
This indicates the formation of intermediate free radicals
during oxidation.
Product analysis and stoichiometry
RESULTS
A 20 cm3 volume of 1 M hexacyanoferrate(III) solution
was mixed with 20 cm3 of a 0.1 M solution of 4-oxo-4-
arylbutanoic acid in sodium hydroxide and the reaction
was allowed to go to completion in a thermostat at 40 ꢀC
for 3–4 h. The solution was then cooled and acidified. A
precipitate was formed along with effervescence due to
the evolution of carbon dioxide. The precipitate was
extracted into diethyl ether, which on evaporation gave
benzoic acid, identified by TLC and chemical analysis. It
is well known that 4-oxo-acids on oxidation give ꢁ,ꢂ-
unsaturated ketones,17,18 but when the oxo acid is un-
substituted in the ꢁ- or ꢂ-position, ꢁ,ꢂ-unsaturated
ketones are not isolated.19 The possibility of phenyl vinyl
ketone being formed in our experiments cannot be ruled
out. Although it may be formed, it is highly susceptible to
nucleophilic attack in alkaline medium.20 To obtain a
The kinetics of oxidation of unsubstituted and substituted
4-oxo-4-arylbutanoic acids by hexacyanoferrate(III) in
aqueous alkaline solution are found to be first order with
respect to the substrate, oxidant and OHꢁ ions. At fixed
[substrate] and [OHꢁ] which are always in excess over
the [oxidant], the plots of log[oxidant] against time are
linear for all the substrates for at least two half-lives of
the oxidant. The pseudo-first-order rate constants calcu-
lated from the first-order plots are given in Table 1.
Although a decrease in the rate constant occurs with
increase in the initial concentration of hexacyanoferra-
te(III) [Table 1(a)], the reaction shows no deviation in the
first-order dependence on the oxidant at any concentra-
tion. Such a decrease in the rate constant has also been
reported in the oxidation of other oxo compounds.3–6
Table 1. Pseudo-first order rate constants for the oxidation of 4-oxo-4-arylbutanoic acids: effect of variation of (a) [oxidant]0,
(b) [substrate] and (c) [OHꢁ]a
(a) [Substrate] ¼ 2 ꢂ 10ꢁꢁ23mol dmꢁ3; [OHꢁ] ¼ 1.5 ꢂ 10ꢁ2 mol dmꢁ3; I ¼ 0.118
103 [oxidant] 0 (mol dm
)
2.0
1.44
3.0
1.40
4.0
1.30
5.5
1.22
8.0
1.09
10.0
0.96
12.0
0.82
103k (sꢁ1
)
(b) [Oxidant] ¼ 2 ꢂ 10ꢁ3 mol dmꢁ3; [OHꢁ] ¼ 5.4 ꢂ 10ꢁ2 mol dmꢁ3; I ¼ 0.082
103 [suꢁbs1trate] (mol dmꢁ3
)
6.0
1.14
8.0
1.50
10.0
1.90
12.0
2.20
16.0
3.00
103k (s
)
(c) [Oxidant] ¼ 2 ꢂ 10ꢁ3 mol dmꢁ3; [substrate] ¼ 1 ꢂ 10ꢁ2 mol dmꢁ3; I ¼ 0.108
102 [OHꢁ] (mol dmꢁ3
)
4.0
1.51
5.0
1.92
6.0
2.38
7.0
2.64
8.0
3.01
103k (sꢁ1
)
a [OHꢁ] refers to the concentration of free hydroxide ions left after neutralization of the carboxyl group in the substrate.
Copyright # 2005 John Wiley & Sons, Ltd. J. Phys. Org. Chem. 2005; 18: 1042–1049

1044
F. J. MARIA PUSHPARAJ ET AL.
Table 2. Effect of varying the ionic strength in the oxidation of 4-oxo-4-phenylbutanoic acids by hexacyanoferrate(III) at 313 K
[Substrate] ¼
1.0 ꢂ 10ꢁ2 mol dmꢁ3
;
[OHꢁ] ¼ 1.0 ꢂ 10ꢁ2 mol dmꢁ3
;
[oxidant] ¼ 2.0 ꢂ 10ꢁ3 mol dmꢁ3
I
0.038
1.84
0.050
2.42
0.060
2.54
0.074
3.58
0.098
4.32
0.134
5.12
0.158
5.53
0.188
6.89
0.218
8.21
104k (sꢁ1
)
Further, the reaction is found to be specifically hydroxide
catalysed since other bases such as acetate ions fail to
bring about the reaction.
Effect of adding hexacyanoferrate(II)
The addition of hexacyanoferrate(II) decreases the rate of
reaction (Table 4). The dependence of the rate on hex-
acyanoferrate(II) is in the inverse order.
Effect of ionic strength
An increase in ionic strength results in an increase in the
½
rate constant (Table 2) and a plot of log kobs against I
(Fig. 1) gives a limiting slope of 5.0.
Substituent effects and evaluation
of Arrhenius parameters
The oxidations of different phenyl-substituted oxo acids
(ArCOCH2CH2CO2H, where Ar ¼ 4-methoxy, 4-ethoxy,
4-methyl, 3,4-dimethyl, 2,4-dimethyl, 2,4,4-trimethyl, 4-
chloro, 4-bromo and 3-nitro substituted phenyl) were
studied. The study was also extended to 4-oxo-4-(1-
naphthyl)- and 4-oxo-4-(2-naphthyl)butanoic acids. In
all these cases, the reaction orders are the same, namely
first order with respect to oxidant, substrate and hydro-
xide. Electron-withdrawing substituents in the phenyl
ring enhance the rate of oxidation and electron-releasing
substituents decrease it. The Hammett plots (Fig. 2) are
linear with good correlation coefficients at four different
Effect of solvents
A decrease in the relative permittivity ("r) of the medium,
effected by adding methanol, decreases the rate of oxida-
tion considerably. However, on addition of dimethyl
sulfoxide the rate of oxidation decreases by up to 40%,
then increases (Table 3).
Figure 1. Effect of varying ionic strength
Figure 2. Hammett plots
Table 3. Effect of addition of methanol and dimethyl sulf-
oxide on the oxidation of 4-oxo-4-phenylbutanoic acids by
hexacyanoferrate(III) at 308 K
Table 4. Effect of hexacyanoferrate(II) on the oxidation 4-
oxo-4-phenylbutanoic acids by alkaline hexacyanoferrate(III)
at 308 K
[Subꢁstrate] ¼ 1.0 ꢂ 10ꢁ2 mol dmꢁ3
;
[OH ] ¼ 2.0 ꢂ 10ꢁ2 mol dmꢁ3; [oxidant] ¼ 2 ꢂ 10ꢁ3 mol dmꢁ3
[Substrate] ¼ 1.0 ꢂ 10ꢁ2 mol dmꢁ3
;
[OHꢁ] ¼ 2.0 ꢂ 10ꢁ2 mol
dmꢁ3; [oxidant] ¼ 2.0 ꢂ 10ꢁ3 mol dmꢁ3; I ¼ 0.042
Methanol (%, v/v)
104k (sꢁ1
DMSOꢁ(1%, v/v)
104k (s
0
5
10
1.46
40
1.90
15
1.26
50
2.28
20
0.78
60
3.0
)
3.42
10
5.29
2.8
20
3.06
4ꢁ
103 [Fe(CN)6
104k (sꢁ1
0
6.2
1.0
4.78
2.0
4.5
3.0
3.97
)
)
Copyright # 2005 John Wiley & Sons, Ltd.
J. Phys. Org. Chem. 2005; 18: 1042–1049

ALKALINE HEXACYANOFERRATE(III) OXIDATION OF 4-OXO ACIDS
1045
Table 5. Substituent effect and activation parameters in the oxidation of ArCO(CH2)2CO2H by alkaline hexacyanoferrate(III)
[Substrate] ¼ 1.0 ꢂ 10ꢁ2 mol dmꢁ3; [OHꢁ] ¼ 2.0 ꢂ 10ꢁ2 mol dmꢁ3; [oxidant] ¼ 2 ꢂ 10ꢁ3 mol dmꢁ3; I ¼ 0.042
104k (sꢁ1
)
Ea
ÁHzꢁ1
ÁSz
ÁGzꢁ1
(kJ molꢁ1
)
(kJ mol
)
(J molꢁ1 Kꢁ1
)
(kJ mol
)
Ar
ꢃ
299 K
308 K
313 K
318 K
Phenyl
2-Naphthyl
4-Bromophenyl
1-Naphthyl
3-Nitrophenyl
4-Chlorophenyl
4-Ethoxyphenyl
4-Methoxyphenyl
4-Methylphenyl
2,4-Dimethylphenyl
3,4-Dimethylphenyl
2,4,6-Trimethylphenyl
Reaction constant (ꢀ)
0
0.17
2.56
5.56
6.68
3.42
6.00
8.04
13.8
35.16
8.06
1.80
1.70
2.18
2.88
1.86
10.10
1.36
4.20
6.92
8.90
14.92
34.76
8.96
2.24
2.20
2.24
3.84
2.56
10.70
1.25
5.06
7.32
10.8
15.4
32.48
11.0
2.52
2.28
3.02
4.44
2.88
12.6
1.19
29.0
21.7
19.8
13.9
ꢁ4.2
20.7
36.0
33.2
29.3
35.2
39.2
20.0
26.5
19.1
17.2
11.3
ꢁ6.7
18.1
33.4
30.6
26.7
32.6
36.6
17.4
ꢁ155.8
ꢁ174.9
ꢁ178.5
ꢁ193.2
ꢁ244.0
ꢁ175.7
ꢁ138.6
ꢁ148.1
ꢁ158.3
ꢁ137.3
ꢁ127.9
ꢁ175.6
74.5
73.0
72.2
70.8
68.4
72.2
76.1
76.2
75.5
74.9
76.0
71.5
þ0.23
þ0.40 11.2
þ0.71 36.14
þ0.23
ꢁ0.25
ꢁ0.27
ꢁ0.17
—
6.66
1.10
1.08
1.46
1.98
1.18
7.76
1.57
ꢁ0.24
—
temperatures: 299 K, 308 K, 313 K and 318 K (com-
pounds with substituents in the 2- or 2,6-position of the
benzene ring are not included for obvious reasons). The
reaction constants (ꢀ) are positive and decrease with
increasing temperature.
The Arrhenius relationship is observed in all cases but
with poorer correlation coefficients than those of the
Hammett plots. The reactions are characterized by low
values of activation energy and large negative values of
entropy of activation. Table 5 gives the values of the rate
constants at four temperatures and also the Arrhenius
parameters.
k2½A2ꢁꢃ½Fꢃ
½RAꢃ ¼
ð2Þ
kꢁ2½F1ꢃ þ k3
½A2ꢁꢃ½H2Oꢃ
k1½HAꢁꢃ½OHꢁꢃ
½H2Oꢃ
½k1ꢃ ¼
or ½A2ꢁꢃ ¼
ð3Þ
ð4Þ
½HAꢁꢃ½OHꢁꢃ
k1k2½HAꢁꢃ½OHꢁꢃ½Fꢃ
À
Á
½RAꢃ ¼
k2½F1ꢃ þ k3 ½H2Oꢃ
It is interesting to note that the experimentally obtained
energy of activation is negative for the oxidation of 4-
oxo-4-(3-nitrophenyl)butanoic acid.
rate ¼ k3½RAꢃ½H2Oꢃ (assuming slow step)
k1k2k3½HAꢁꢃ½OHꢁꢃ½Fꢃ
rate ¼
ð5Þ
ð6Þ
kꢁ2½F1ꢃ þ k3
DISCUSSION
Mechanism
k1k2k3½HAꢁꢃ½Fꢃ½OHꢁꢃ
rate ¼
kꢁ2½F1ꢃ
Based on the above results and in analogy with the
alkaline hexacyanoferrate(III) oxidation of other oxo
compounds, the mechanism shown in Scheme 1 may be
proposed, where K (k1/kꢁ1) is the equilibrium constant
for enolate anion (A2ꢁ) formation.
Equation (5) shows a negative fractional order in
ferrate(II) and Eqn (6) supports a negative fractional
order. This explains the observed first-order dependence
on the oxidant, substrate and OHꢁ. The experimental rate
constant when divided by the concentrations of the oxo
acid and OHꢁ gives a value of k2K which is the product of
the rate constant for the electron transfer step and of the
equilibrium constant for the proton abstraction. The
formation of enolate anion is a prerequisite for the
oxidation,21 as benzaldehyde, which cannot enolize, is
not oxidized by alkaline hexacyanoferrate(III).
Steady-state equation:
d=dtðRAÞ ¼ k2½A2ꢁꢃ½Fꢃ ꢁ kꢁ2½RAꢃ½F1ꢃ ꢁ k3½RAꢃ ¼ 0
ð1Þ
where RA ¼ radical anion
À
Á
From the above mechanism, the experimental results
can be explained.
½RAꢃ kꢁ2½F1ꢃ þ k3 ¼ k2½A2ꢁꢃ½Fꢃ
Copyright # 2005 John Wiley & Sons, Ltd.
J. Phys. Org. Chem. 2005; 18: 1042–1049

1046
F. J. MARIA PUSHPARAJ ET AL.
Scheme 1
Decrease of rate constant with increasing
[oxidant]0
values obtained were 2.73, 9.1 ꢂ 10ꢁ4 and 2.7 ꢂ
10ꢁ4 molꢁ1 dm3 minꢁ1. From earlier investigations on
the oxidation of oxo compounds,3,5,6 k1 was calculated
in this way for ethyl methyl ketone, formadehyde and
pyruvic acid. The value thus obtained for ethyl methyl
ketone agrees well with that found in the literature.23,24
The k1 values obtained for the oxidation of 4-oxo acids
are given in Table 6.
The Hammett plot drawn with these k1 values gives a ꢀ
value of 0.69, which is comparable to the value of 0.88
reported for the alkaline hydrolysis of esters of these oxo
acids.24
The reason for the decrease in rate constant with increas-
ing [hexacyanoferrate(III)]0 may be found by considering
the sequence of reaction steps. In the rate equation
obtained from the proposed mechanism, it is assumed
that kꢁ1[H2O] is much greater than k2 [F]. At higher
concentrations of hexacyanoferrate(III), k2[F] may reach
a value comparable to k1[H2O] and may no longer be
ignored.
From equation (5),
k1k2k3½HAꢁꢃ½OHꢁꢃ½Fꢃ
kobs
¼
kꢁ2½F1ꢃ þ k3
Effect of ionic strength
The large increase in the rate constants with increasing
ionic strength of the medium clearly points to a reaction
and
1
k1k2k3½HAꢁꢃ½OHꢁꢃ½Fꢃ
kꢁ2½F1ꢃ þ k3
¼
Table 6. Values of the rate cꢁonstants for proton abstraction
kobs
from XC6H4COCH2CH2 CO2 by OHꢁ ions at 308 K
Therefore, a plot of 1/kobs versus [F] must give a straight
line from the slope of which k1 can be calculated by
following the method suggested by Maskill.22 The k1
X
4-OCH3
0.0255
H
4-Cl
3-NO2
0.115
k1 molꢁ1 dmꢁ3 sꢁ1
0.0328 0.0592
Copyright # 2005 John Wiley & Sons, Ltd.
J. Phys. Org. Chem. 2005; 18: 1042–1049

ALKALINE HEXACYANOFERRATE(III) OXIDATION OF 4-OXO ACIDS
1047
between similarly charged ions. The observed rate con-
stant is the product of the equilibrium constant of the
proton transfer equilibrium and the rate constant of the
oxidation step (Kk2). The Bjerrum–Brønsted equation
relates k2 to ionic strength as
sium ions constant while varying the ionic strength
(Table 7).
It can be seen that if the potassium ion concentration is
kept constant, the rate constants remain unaltered in spite
of the change in ionic strength. The application of the
Olson–Simonson rule underscores the participation of
negatively charged ions in the rate-determining step.
log k2 ¼ constant ꢁ 2QZA ZB I 0:5
Consideration of the charges of enolate and hexacyano-
ferrate(III) ions results in the equation
Arrhenius parameters
In Equation (3), K/[H2O] can be equated to Ka/Kw, where
Ka is the ionization constant of the CH2 adjacent to the
oxo group. Being the ionization constant of a weak acid,
Ka does not vary much with temperature,27,28 whereas the
ionic product of water, Kw, increases with temperature in
the temperature range studied.29 Since the equilibrium
constant for the proton transfer is related to Ka/Kw, its
value decreases with temperature. The observed rate
constant in the oxidation of 4-oxo acids by alkaline
hexacyanoferrate(III) is the product of this equilibrium
constant and the rate constant of the oxidation step. The
former decreases and the latter increases with tempera-
ture and the net result is that observed rate constants vary
by only a small amount with temperature. This is re-
flected in the low values of the ‘activation energy’, which
is in fact the sum of the enthalpy change for the proton
transfer equilibrium (which is negative) and the activa-
tion energy of the electron abstraction step. The reason
for the decrease in the observed rate constant with
temperature in the oxidation of 4-oxo-4-(3-nitrophenyl)-
butanoic acid is presumably due to the fact that the
enthalpy change for the proton transfer is greater in
magnitude than the activation energy of the oxidation
step.
log k2 ¼ constant ꢁ 12QI 0:5
The equilibrium constant K and I are related as
ꢁ
ꢁ
log K ¼ constant þ log fH2A ꢁ log fHA ꢁ log fOH
According to the Debye-Hu¨ckel limiting law:
log fi ¼ ꢁQZi2I0:5
Therefore, log K ¼ constant ꢁ 2QI0.5. Combining the
equations for k2 and K,
log k2K ¼ constant þ 10I0:5
At 40 ꢀC, Q has a value of 0.52, so log k2K ¼ constant +
5.2 I0.5
.
The limiting slope obtained from Fig. 1 is 5.0, which is
in agreement with the above equation. The deviation
from linearity at higher ionic strengths is due to failure
of the Bjerrum–Brønsted equation and the Debye-Hu¨ckel
limiting law.
Attempts made to determine the equilibrium constant
for enolate anion formation have not been successful. The
classical method30 of treating the oxo compound with
bromine and titrating the excess bromine met with fail-
ure, as the oxo acid was completely oxidized by bromine
in a very fast step. The e.m.f. method is unreliable in
alkaline media.31 Insufficient data are available for the
determination of the equilibrium constant by the use of
the Marcus equation.32
Olson–Simonson rule26
In a reaction between ions of similar charge, the con-
centration of oppositely charged ions influences the rate
constant more than the ionic strength of the medium. The
application of this rule was attempted in the present
oxidation studies, by keeping the concentration of potas-
Table 7. Application of Olson–Simonson rule
The entropy of activation is negative in all the cases.
This is due to the increased electrostriction when similar
charges come together in the activated state. In aqueous
solution in a reaction between ions of charge zA and zB,
the entropy of activation has been shown33 to be ap-
proximately equal to ꢁ41zAzB J Kꢁ1 molꢁ1. Therefore, in
the present studies, the entropy of activation for the
oxidation step, involving enolate anion of charge 2ꢁ
and hexacyanoferrate(III) of charge 3ꢁ, should be about
ꢁ246 J Kꢁ1 molꢁ1. It is seen that the observed entropies
of activation are greater than this value, presumably
owing to the positive entropy change in the proton
transfer equilibrium.
[Substrate]0 ¼ 1.0 ꢂ 10ꢁ2 mol dmꢁ3;[OHꢁ] ¼ 2.0 ꢂ 10ꢁ2 mol
dmꢁ3; [oxidant]0 ¼ 3.0 ꢂ 10ꢁ3 mol dmꢁ3; T ¼ 313K
Added salt
[Salt]ꢁ3
(mol dm
[Kþ]ꢁ3
(mol dm )
I
104 kobs
)
None
KCl
K2SO4
KCl
K2SO4
KCl
K2SO4
—
0.012
0.027
0.0
0.038
0.050
0.056
0.098
0.121
0.218
0.308
2.76
3.60
3.64
6.50
6.20
12.3
12.0
0.006
0.06
0.03
0.18
0.09
0.041
0.089
0.089
0.209
0.209
Copyright # 2005 John Wiley & Sons, Ltd.
J. Phys. Org. Chem. 2005; 18: 1042–1049

1048
F. J. MARIA PUSHPARAJ ET AL.
Structure–reactivity correlations
Oxidation reactions usually have a low negative ꢀ value.
In the present studies, the overall rate constant is the
product of an equilibrium constant and a rate constant.
The equilibrium leading to the formation of electron-rich
enolate anion would necessarily have a positive ꢀ and the
subsequent oxidation would have a low negative ꢀ. The
net effect would therefore lead to a positive ꢀ for the
overall reaction. The application of the modified Ham-
mett equation, which incorporates the inductive and
resonance parameters34 of para-substituted compounds,
shows that the inductive effect is 52.4% and the reso-
nance effect is 47.6%.
Figure 3. Arrhenius plots
Ortho effect
Details of electron abstraction
It is seen that whereas a methyl substituent in the para
position of the phenyl ring is generally rate retarding, its
introduction in the ortho position increases the rate
considerably. Methyl groups in the two ortho positions
increase the rate almost to that of the 4-chloro compound.
Hence the reaction is subjected to steric acceleration by
the ortho substituents.
The electron loss from the enolate dianion can occur
either from the oxygen atom of the enolate function or
from the carbon–carbon double bond. Both processes will
lead to the formation of a mesomeric radical:
Oꢁ
C
k
j
ꢁ
ꢁ
ꢁ
---
---
---
---
---
---
---
C6H5---- C ---CH CH2 COO $ C6H6 C --- C H CH2 COO
Isokinetic relationship (IKR)
Generally, one-electron-abstracting agents do not read-
ily attack the olefinic bonds.21 The loss of electron from
the carboxylate moiety of the molecule is less probable,
as the substituent effects would have been negligible. In
the oxidation of phenylacetic acid by Co(III), an electron
is proposed to be removed from the phenyl ring rather
than the carboxyl group, resulting in the aromatic radical
cation and a highly negative ꢀþ is obtained.37 The second
electron removal occurs in a fast step with decarboxyla-
tion occurring at the same time, in accordance with the
Kochi mechanism.38
This explains the formation of CO2 in the oxidation of
4-oxo acids. Owing to the high reactivity of acrylophe-
none in alkaline solution, only benzoic acid is obtained in
the oxidation of 4-oxo-4-phenylbutanoic acid, whereas
chalcone is formed in the oxidation of 4-oxo-2,4-diphe-
nylbutanoic acid.
The Hammett constant (ꢀ) decreases with increase in
temperature, from 1.57 at 299 K to 1.19 at 318 K. The
Hammett lines intersect at a point corresponding to a ꢃiso
value of 0.60 and a kiso value of 3.090 ꢂ 10ꢁ3 sꢁ1 (Fig. 2).
An oxo acid with a substituent having this ꢃ value will be
oxidized by alkaline hexacyanoferrate(III) with the same
velocity at all temperatures. Such an insensitivity of the
reaction rate to temperature is a pointer to the existence of
an isokinetic phenomenon. Introduction of a nitro group
in the phenyl ring which has a higher ꢃ value results in a
decrease in the rate constant with temperature.
A reaction series which exhibits a common point of
intersection in the Hammett plot is expected to have a
common point of intersection in the Arrhenius plot also.35
The Arrhenius plots are shown in Fig. 3.
As the common point of intersection is not easily
visible, the problem is tackled by the statistical method
developed by Exner and Beranek.36 In this method, the
standard deviation S00 of the free regression lines (with-
out the constraint of a common point of intersection) and
the standard deviation S0 of the new constrained lines are
compared. In the present studies, these two values are
approximately equal and so the isokinetic hypothesis
cannot be rejected. From the plots of ÁHz against ÁSz
and ꢀ against 1/T, the Tiso value was found to be 385 K.
It is interesting that although the rate constant is
composite—k2K, involving two reaction steps—an
anti-compensation effect is not observed.
CONCLUSION
From the Hammett and Arrhenius plots associated with
the isokinetic relationship, it is noted that the ꢀ values are
positive and decrease with increase in temperature. A
methyl group in the ortho position accelerates the reac-
tion owing to steric factors, contrary to the expected
electronic effect. The structure–activity correlations are
made on the basis of the results obtained and suggest a
mesomeric radical-mediated mechanism.
Copyright # 2005 John Wiley & Sons, Ltd.
J. Phys. Org. Chem. 2005; 18: 1042–1049

ALKALINE HEXACYANOFERRATE(III) OXIDATION OF 4-OXO ACIDS
1049
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The spectral data were obtained using instruments funded
by DST-FIST. Financial assistance from the UGC SAP
(Phase I)-DRS programme is greatly acknowledged.
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Copyright # 2005 John Wiley & Sons, Ltd.
J. Phys. Org. Chem. 2005; 18: 1042–1049