Aquachlororuthenium(III) catalysis in the oxidation of substituted 4-oxo-4-arylbutanoic acids by bromate in acid medium: A kinetic and mechanistic study and validity of linear free-energy relationships

Article abstract of DOI:10.1007/s11243-011-9523-x

Ru(III) acts an efficient catalyst in the oxidation of substituted 4-oxo-4-arylbutanoic acids (4-oxo acids) by bromate in sulfuric acid medium, giving the corresponding benzoic acids in quantitative yields. The reaction shows first-order dependence in both [bromate] and [H₂SO ₄], and a non-linear dependence on both [oxo acid] and [catalyst]. Changing solvent from H₂O to D₂O increases the rate. The rate is not affected by ionic strength but decreases with increase in dielectric constant of the medium. Electron-releasing substituents in the phenyl ring of the substrate greatly accelerate the rate, whereas the retardation by electron-withdrawing substituents, though perceptible, is small. The linear free-energy relationship is characterized by smooth curves in Hammett plots of log k versus σ; however, linear plots are obtained with excellent correlation coefficients at all the studied temperatures, when Brown's σ⁺ values are used. The reaction constant is negative and decreases with increase in temperature. From the intersection of the lines in the Hammett and Arrhenius plots, the isokinetic relationship is evaluated. A mechanism involving a cyclic oxidant-substrate-catalyst ternary complex is proposed, in which both C-C bond-breaking and C-O bond formation are involved, and the oxidation state of Ru(III) remains unchanged. A rate law explaining all the kinetic results has been derived and verified. The reaction is an example of neighboring group participation in intramolecular catalysis and is potentially useful for the synthesis of substituted benzoic acids.

Full text of DOI:10.1007/s11243-011-9523-x

Transition Met Chem (2011) 36:707–719
DOI 10.1007/s11243-011-9523-x
Aquachlororuthenium(III) catalysis in the oxidation of substituted
4-oxo-4-arylbutanoic acids by bromate in acid medium: a kinetic
and mechanistic study and validity of linear free-energy
relationships
•
Padma Sunitha Manjari Cherkupally Sanjeeva Reddy
Received: 27 May 2011 / Accepted: 25 July 2011 / Published online: 12 August 2011
Ó Springer Science+Business Media B.V. 2011
Abstract Ru(III) acts an efficient catalyst in the oxida-
tion of substituted 4-oxo-4-arylbutanoic acids (4-oxo acids)
by bromate in sulfuric acid medium, giving the corre-
sponding benzoic acids in quantitative yields. The reaction
shows first-order dependence in both [bromate] and
[H2SO4], and a non-linear dependence on both [oxo acid]
and [catalyst]. Changing solvent from H2O to D2O
increases the rate. The rate is not affected by ionic strength
but decreases with increase in dielectric constant of the
medium. Electron-releasing substituents in the phenyl ring
of the substrate greatly accelerate the rate, whereas the
retardation by electron-withdrawing substituents, though
perceptible, is small. The linear free-energy relationship is
characterized by smooth curves in Hammett plots of log
k versus r; however, linear plots are obtained with excel-
lent correlation coefficients at all the studied temperatures,
when Brown’s r? values are used. The reaction constant is
negative and decreases with increase in temperature. From
the intersection of the lines in the Hammett and Arrhenius
plots, the isokinetic relationship is evaluated. A mechanism
Introduction
The carbonyl and carboxy groups of 4-oxo-4-phenylbuta-
noic acid and its phenyl substituted compounds (commonly
known as 4-oxo acids) are separated by two carbons;
hence, the two functional groups are essentially indepen-
dent. However, an intramolecular mechanism in the cata-
lytic iodination of 4-oxo acid [1] and neighboring group
participation in the hydrolysis of its ethyl ester [2, 3] has
been reported. In acidic medium 4-oxo acids undergo
enolization [4], and on selective oxidation, they result in
malonic and substituted benzoic acids, which find wide-
spread applications in synthetic organic chemistry. Hence,
4-oxo acids are attractive substrates in terms of their
mechanistic aspects.
Sodium and potassium bromate are stable solids and can
be easily handled when compared with liquid bromine
and hypobromous acid solutions. The product of bromate
oxidation is bromide, which can be safely recycled, thus
making bromate oxidations environmentally benign com-
pared with metal ion oxidations. Though bromate itself is
a strong oxidizing agent, having a redox potential of
1.45 V [5], the oxidations generally require a catalyst.
Ru(III) has been used as a homogeneous catalyst in many
redox reactions and its versatility as a catalyst is revealed
by oxidation studies of various organic substrates involving
different oxidants [6–10]. The mechanism of catalysis is
often quite complicated due to the formation of different
intermediate complexes, free radicals and oxidation states
of Ru.
involving
a cyclic oxidant–substrate–catalyst ternary
complex is proposed, in which both C–C bond-breaking
and C–O bond formation are involved, and the oxidation
state of Ru(III) remains unchanged. A rate law explaining
all the kinetic results has been derived and verified. The
reaction is an example of neighboring group participation
in intramolecular catalysis and is potentially useful for the
synthesis of substituted benzoic acids.
Our preliminary experimental results indicated that the
reaction of 4-oxo acids with bromate in acid medium
without a catalyst is sluggish, but accelerated by a small
amount of RuCl3 catalyst. Therefore, in order to explore
the mechanism of oxo acid–bromate reactions and also to
P. S. Manjari ꢀ C. S. Reddy (&)
Catalysis Laboratory, Department of Chemistry, University
College, Kakatiya University, Warangal 506009, India
e-mail: chsrkuc@yahoo.co.in
123

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Transition Met Chem (2011) 36:707–719
study the catalytic action of Ru(III), we chose RuCl3 as a
catalyst. While studying the present reaction in the pres-
ence of mercury(II), a bromo-complexing metal ion, some
interesting results were obtained. Ru(III)-catalyzed bro-
mate oxidation of malonic acid (MA) and methyl malonic
acid (MMA) exhibit an induction period [11], whose
duration is inversely proportional to [substrate], [catalyst]
and [acid], and independent of [bromate]. After induction,
the reaction is first order in both [bromate] and [catalyst]
but shows less than unity order in [substrate] and an inverse
fractional order in [acid]. However, the present reaction
does not show any induction period over a wide concen-
tration range of bromate, Ru(III), oxo acid and H2SO4. The
kinetic behavior of the present reactants and catalyst is
entirely different from the earlier reports of Ru(III)-cata-
lyzed oxidation reactions [6–11]. The mechanism neither
resembles catalysis by bromide anions [12] nor that of
Ru(III)-catalyzed oxidation of organic substrates with the
oxidants [6–11].
identification of organic intermediates in the reaction were
performed using HPLC. The experiments were performed
with a Shimadzu instrument using an ion-exchange column
at 45 °C and a UV detector working at 220 nm. The
intermediates and products were identified from their
retention time (tr). A Shimadzu multipurpose recording
double beam UV–visible spectrophotometer equipped with
a temperature controller was used for absorption studies.
Kinetic measurements
All the kinetic measurements were carried out in black-
coated vessels at constant temperature ( 0.1 °C) under
pseudo-first-order conditions with [oxo acid] ꢁ [bromate].
The reaction was initiated by the rapid addition of tem-
perature-equilibrated oxidant of the required concentration
to reaction mixtures containing the required amounts of
substrate, catalyst, Hg(OAc)2, H2SO4, AcOH and water, in
glass-stoppered Pyrex boiling tubes that were thermally
equilibrated for 1 h. The progress of the reaction was
monitored by iodometric determination of unconsumed
[bromate] in known aliquots of the reaction mixtures at
different time intervals. Before adopting the iodometric
method, it was ensured that the presence of oxo acid in the
quenching solution of KI did not change the bromate titer
value. The course of the reaction was studied for at least
three half-lives. The rate constants (k, s-1) were deter-
mined from the pseudo-first-order plots of log [oxidant]
against time. The pseudo-first-order plots were linear
(R2 C 0.99) for more than 80% completion of the reaction,
and the k (s-1) values were reproducible to within 5%.
Rate constants did not alter in a nitrogen atmosphere,
and all the rate constants reported in this paper were
obtained without nitrogen. Freshly prepared solutions of
oxo acids in purified acetic acid were used throughout. The
uncatalyzed reaction is very slow under these conditions
and is first order in both [bromate] and [oxo acid], and
second order in [H?]. However, the reaction rate is
appreciably faster (40–80 times) in the presence of a
minute quantity (10-5 mol dm-3) of RuCl3. All the rate
constants for the catalyzed reactions were obtained as:
kcatalyzed = koverall – kuncatalyzed and represented as k (s-1).
In the present study, we describe the kinetics and
mechanism of oxidation of substituted 4-oxo-4-arylbuta-
noic acids by bromate in sulphuric acid medium containing
mercury(II) and Ru(III), including the active species of the
substrate, catalyst and oxidant, and the oxidation products.
We have also studied the catalytic efficiency of Ru(III) and
evaluated the related kinetic and thermodynamic parame-
ters of the reaction, along with the linear free-energy and
isokinetic relationships.
Experimental
All the oxo acids were obtained from Aldrich and recrys-
tallized twice from water before use. The purity of sub-
strates was checked by their melting points, UV, IR and
NMR spectra. KBrO3 (Reidal), H2SO4 and Hg(OAc)2
(Merck) were of analytical reagent grade and used as
received. D2O (99.4% pure) was obtained from the Baba
Atomic Research Centre, Mumbai, India. Acetic acid
(BDH) was purified by refluxing with chromic acid and
acetic anhydride for 6 h and then distilled. Ruthenium
trichloride hydrate (Johnson Matthey) was purified by
evaporation with conc. HCl. A stock solution of Ru(III)
was prepared by dissolving a weighed quantity of RuCl3 in
dilute HCl of known strength. Mercury was added to this
solution, to detect the formation of any Ru(IV) during its
preparation, and the solution was kept for 1 day. The
obtained solution was assayed by EDTA titration [13] and
stored in a glass-stoppered flask, painted black, at *5 °C.
Dilute solutions of Ru(III) were made from the stock
solution as required. Other solutions were prepared with
either doubly distilled water or purified acetic acid and
were standardized by known methods. Separation and
Results
Characterization of the reaction kinetics
Molecular bromine production anticipated from the bro-
mate–bromide reaction and its subsequent reaction with the
oxo acid was eliminated by the addition of the bromo-
complexing agent Hg(II) [14]. The added mercury acetate
had no effect on the rate of reaction over a wide
123

Transition Met Chem (2011) 36:707–719
709
concentration range (0.001–0.01 mol dm-3). An optimum
concentration (0.005 mol dm-3) of mercury acetate was
employed throughout; hence, the reported rate constants
and thermodynamic parameters pertain to pure bromate
(uncontaminated with Br2) oxidation only.
obtained pseudo-first-order rate constant (k, s-1), evaluated
from the slopes of such plots, remained unchanged
(Table 1), confirming the first-order dependence of the rate
on [bromate]init
.
The effect of [oxo acid] on the reaction rate was studied
in the concentration range of 0.005–0.03 mol dm-3, at
constant concentrations of oxidant, acid and catalyst
(Table 1). The plots of log k versus log [oxo acid] were
linear (R2 C 0.99) with slope values less than one
(0.50–0.54). The plots of k against [oxo acid] were non-
linear, but the double reciprocal plots, 1/k against 1/[oxo
At fixed concentration of other reactants and when [oxo
acid] is in tenfold excess over [bromate], the kinetic studies
were carried out at different [bromate]init in the range of
5.0–20.0 9 10-4 mol dm-3. Linear plots were obtained
(R2 C 0.99) for log [bromate]t versus time at all [bro-
mate]init for more than three half-lives of the reaction. The
Table 1 Dependence of rate on the factors influencing the Ru(III) catalyzed oxidation of 4-oxo acids by bromate in acid medium at 313 K
Non-variable constituents (mol dm-3
)
Variable constituent
(mol dm-3
104 9 k (s-1
)
)
–H
p-OCH3
p-CH3
p-Cl
m-NO2
[Bromate]
[Oxo acid] = 1.0 9 10-2
[H2SO4] = 1.0
[Ru(III)] = 3.3 9 10-5
AcOH–H2O = 1:1 (% v/v)
[Hg(OAc)2] = 5.0 9 10-3
5.0 9 10-4
7.5 9 10-4
10.0 9 10-4
15.0 9 10-4
20.0 9 10-4
[Oxo acid]
5.0 9 10-3
7.5 9 10-3
10.0 9 10-3
15.0 9 10-3
20.0 9 10-3
25.0 9 10-3
30.0 9 10-3
[Ru(III)]
12.79
12.74
12.79
12.78
12.81
115.15
115.09
115.15
115.10
115.14
28.78
28.76
28.78
28.77
28.78
7.89
1.75
1.75
1.75
1.74
1.76
7.95
7.99
7.99
7.94
[Bromate] = 1.0 9 10-3
[H2SO4] = 1.0
[Ru(III)] = 3.3 9 10-5
AcOH–H2O = 1:1 (% v/v)
[Hg(OAc)2] = 5.0 9 10-3
9.59
10.99
12.79
15.32
18.19
19.87
22.90
79.94
96.60
115.15
143.91
169.82
192.75
–
19.67
23.71
28.78
34.76
41.22
45.18
51.16
5.53
6.91
1.07
1.46
1.75
2.18
2.55
2.87
3.18
7.99
9.54
11.51
12.69
13.90
[Bromate] = 1.0 9 10-3
[Oxo acid] = 1.0 9 10-2
[H2SO4] = 1.0
1.65 9 10-5
3.33 9 10-5
6.66 9 10-5
9.99 9 10-5
13.32 9 10-5
[H2SO4]
8.47
12.79
17.75
20.98
25.58
64.51
115.15
161.29
191.91
227.72
20.89
28.78
43.97
52.48
57.57
5.24
7.99
1.27
1.75
2.50
2.87
3.46
11.51
13.90
15.98
AcOH–H2O = 1:1 (% v/v)
[Hg(OAc)2] = 5.0 9 10-3
[Bromate] = 1.0 9 10-3
[Oxo acid] = 1.0 9 10-2
[Ru(III)] = 3.3 9 10-5
AcOH–H2O = 1:1 (% v/v)
[Hg(OAc)2] = 5.0 9 10-3
0.25
3.22
6.39
28.08
53.71
9.59
19.19
24.21
28.78
46.06
57.57
2.36
4.39
6.30
7.99
9.59
11.51
0.43
0.92
1.44
1.75
2.42
2.87
0.50
0.75
9.59
83.17
1.00
12.79
15.89
19.25
115.15
140.63
172.42
1.25
1.50
AcOH–H2O (% v/v)
30–70 (53.18)a
40–60 (46.48)
50–50 (39.78)
60–40 (33.08)
70–30 (26.08)
[Bromate] = 1.0 9 10-3
[Oxo acid] = 1.0 9 10-2
[Ru(III)] = 3.3 9 10-5
[H2SO4] = 1.0
7.67
9.59
38.38
51.15
115.15
307.04
–
16.79
23.02
28.78
43.17
76.76
4.79
6.39
0.72
1.53
1.75
2.68
3.83
12.79
19.19
33.58
7.99
9.59
[Hg(OAc)2] = 5.0 9 10-3
13.43
Reported rate constants are the mean of duplicate experiments
a
Parentheses values indicate the dielectric constant of the medium
123

710
Transition Met Chem (2011) 36:707–719
Fig. 1 Plot of 1/k against
1/[oxo acid] at 313 K. Reaction
conditions as in Table 1.
A p-OCH3, B p-CH3, C –H,
D p-Cl, E m-NO2
20
15
10
5
D
C
E
k
k
B
0
0
50
100
150
200
250
1 / [ 4-oxo acid] ( dm3 mol –1
)
1 / [ 4-oxo acid] ( dm3 mol –1
)
Fig. 2 Plot of 1/k against
1/[Ru(III)] at 313 K. Reaction
conditions as in Table 1.
A p-OCH3, B p-CH3, C –H,
D p-Cl, E m-NO2
E
C
k
k
B
A
1 / [ Ru(III) ] 10 –4 ( dm3 mol –1
)
1 / [ Ru(III) ] 10 –4 ( dm3 mol –1
)
acid], were linear with varying intercepts and slopes
(Fig. 1), confirming that the order in [oxo acid] varies from
one to zero. Under similar experimental conditions, the
uncatalyzed reaction, though very slow, was first order with
respect to [oxo acid].
to be unity. The plots of 1/k versus 1/[H2SO4] were found
to be linear passing through the origin (Fig. 3). Under
similar conditions, the uncatalyzed reaction was second
order with respect to [H2SO4].
The effect of ionic strength was studied using NaClO4,
with other experimental conditions held constant. There
was no significant effect of ionic strength on the reaction
rate; hence, ionic strength of the medium was not fixed at
any constant value in the other experiments.
The effect of [Ru(III)] was studied between the con-
centration range of 1.65 9 10-5 and 13.32 9 10-5
mol dm-3 (Table 1). The order in [catalyst] was found to
be 0.50 0.05 as determined from the slope of log
k against log [catalyst] plots, and the plots of k against
[catalyst]0.5 were linear. Further, plots of 1/k versus
1/[Ru(III)] (Fig. 2) were linear with varying intercepts on
the 1/k axis for the studied oxo acids.
The dielectric constant of the medium was varied using
different proportions of acetic acid and water (30–70% v/v).
The D values were calculated from the equation
D ¼ ðDW VW Þ þ ðDAVAÞ
The effect of [H2SO4] on the reaction rate was studied
in order to establish the active species of reactants present
in the solution. At fixed concentrations of substrate, bro-
mate, catalyst and other conditions remaining constant, the
reaction rate increased linearly (Table 1) with increase in
[H2SO4] and the order with respect to [H2SO4] was found
where DW and DA are the dielectric constants of pure water
and acetic acid, respectively, and VW and VA are the volume
fractions of components water and acetic acid, respectively
in the total mixture. The rate increased with increasing
acetic acid content (Table 1). Plots of log k versus 1/D were
123

Transition Met Chem (2011) 36:707–719
711
Fig. 3 Plot between 1/k and
1/[H2SO4] at 313 K. Reaction
conditions as in Table 1.
A p-OCH3, B p-CH3, C –H,
D p-Cl, E m -NO2
k
k
1 / [H2SO4] (dm3 mol–1)
1 / [H2SO4] (dm3 mol–1)
linear with positive slopes. Blank experiments showed that
acetic acid was not oxidized by bromate under these con-
ditions. The observed results therefore truly represent the
rates under different dielectric medium of the solution.
isotope effect suggests a proton catalyzed oxidation
reaction.
The oxidation reactions were studied in the temperature
range of 298–323 K keeping all constituents of the solution
constant. The experimental rate constant k (which is a
composite of the enolization, complexation and oxidation
steps) was determined at all the temperatures for all the
studied oxo acids (Table 3). Activation energy (Ea) was
calculated from the least-square slopes of linear Arrhenius
plots of log k versus 1/T (Fig. 4a; R2 C 0.99). The related
thermodynamic parameters, i.e., the enthalpy of activation
(DH#), entropy of activation (DS#) and free energy of acti-
vation (DG#) were calculated using the appropriate equa-
tions and are presented in Table 3. The plot of DH# against
DS# was linear (Fig. 4b), and the isokinetic temperature (b)
obtained is 370.3 K. The value of b calculated from the
Exner’s plot (Fig. 4c) of log k323 K against log k313 K is
370.2 K, which is in good agreement with the value obtained
from the DH# against DS# plot. The value of b is higher than
the experimental temperature range (298–323 K), indicating
that the reaction is enthalpy controlled.
Effects of conditions and substituents
The reactions were neither accelerated nor retarded by
added acrylonitrile (0.1–1.0 mol dm-3) and also no pre-
cipitate due to polymerization of acrylonitrile was
observed, suggesting the absence of any free radicals in the
reaction. To further confirm the absence of free radicals,
the reaction was carried out in the presence of
0.05 mol dm-3 2,6-di-t-butyl-4-methylphenol (BHT). The
unreacted BHT was recovered almost quantitatively at the
end of the experiment.
Rate studies in D2O medium revealed that the reaction
rates were faster in D2O than in H2O and k(H2O)/
k(D2O) = 0.55 0.01 at all the studied temperatures for
all the studied oxo acids (Table 2). Since D3O? is a
stronger acid than H3O? [15], the value of the solvent
The substituent effect on the rate of oxidation was
studied using different phenyl substituted 4-oxo acids
(ArCOCH2CH2COOH, where Ar = 4-methoxy-, 4-methyl-,
4-chloro- or 3-nitro-substituted phenyl) at different tempera-
tures (Table 3). Structural changes in the phenyl moiety result
in varied reactivity; while electron-releasing substituents
greatly accelerate the rate of oxidation, the retardation by
electron-withdrawing substituents though perceptible was
small. Thelinearfree-energyrelationshipis characterizedbya
smooth curve in the Hammett plot of log k versus r at all the
studied temperatures (Fig. 5a). However, linear Hammett
plots were obtained, with excellent correlation coefficients at
all the studied temperatures (Fig. 5b; R2 C 0.998), when
exalted substituent constants (Brown and Okamoto’s r?)
values [16] were used. The reaction constant is negative and
Table 2 Solvent isotope effect on the rate of Ru(III)-catalyzed bro-
mate oxidation of 4-oxo acids at 313 K
4-oxo
acid
k (H2O) 9 104
(s-1
k (D2O) 9 104
(s-1
k (H2O)/
k (D2O)
)
)
–H
12.79
28.78
7.99
23.02
51.22
14.53
3.24
0.55
0.56
0.55
0.54
p-methyl
p-chloro
m-nitro
1.75
Solution conditions: [Oxo acid] = 1.00 9 10-2 mol dm-3, [bro-
mate] = 1.00 9 10-3 mol dm-3, [Ru(III)] = 3.3 9 10-5 mol dm-3
,
[H2SO4] = 1.00 mol dm-3, [Hg (OAc)2] = 5. 00 9 10-3 mol dm-3
and AcOH–H2O = 1:1% (v/v)
Reported rate constants are the mean of duplicate experiments
123

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Transition Met Chem (2011) 36:707–719
Table 3 Temperature dependence and activation parameters for the Ru(III)-catalyzed and uncatalyzed oxidation of 4-oxo acids by bromate in
acid medium
4-oxo
acid
105 9 k (s-1
)
Ea
(kJ mol-1
DH#
DS#
DG#
)
(kJ mol-1
)
(J mol-1 K-1
)
(kJ mol-1
)
298
303
308
313
318
323 K
–H
47.97
(0.39)a
639.22
(14.43)
11.15
(1.13)
19.29
(0.26)
2.98
73.31
(0.70)
767.66
(19.19)
159.93
(1.91)
31.90
(0.39)
5.75
95.95
(1.35)
959.58
(28.78)
223.88
(3.19)
43.61
(0.71)
9.59
127.94
(3.19)
1,151.15
(47.97)
287.87
(6.71)
79.96
191.91
(4.79)
338.22
(9.59)
68.92
66.30
-89.06
(80.26)
-168.5
(-68.32)
-118.8
(-69.47)
-49.9
96.79
(101.46)
38.29
(98.87)
35.69
(73.74)
88.45
p-OCH3
p-CH3
p-Cl
1,279.41
(57.57)
447.72
(11.19)
134.34
(1.72)
1,599.18
(67.17)
639.54
(17.61)
223.80
(3.59)
(47.86)
57.44
(45.26)
54.84
(66.64)
92.04
(76.86)
82.33
(73.96)
79.73
(71.79)
95.36
(1.06)
17.58
(114.85)
105.30
(134.09)
(112.25)
102.70
(131.49)
(9.5)
(70.98)
99.34
m-NO2
223.80
(3.59)
63.90
10.7
(–)
(–)
(0.17)
(0.37)
(1.43)
(66.27)
(76.21)
Solution conditions: [Oxo acid] = 1.00 9 10-2 mol dm-3, [bromate] = 1.00 9 10-3 mol dm-3, mol dm-3, [Ru(III)] = 3.3 9 10-5 mol dm -3
[H2SO4] = 1.00 mol dm-3, [Hg (OAc)2] = 5. 00 9 10-3 mol dm-3 and AcOH–H2O = 1:1% (v/v)
,
The average error in values of Ea, DH#, DS# and DG# are 1.6 kJ mol-1
,
1.6 kJ mol-1
,
3 J mol-1 K-1 and 2 kJ mol-1, respectively.
DH#, DS# and DG# values were obtained at 313 K. The average error in rate constant is 5%
Parentheses values indicate the rate constants (kun, s-1) and activation parameters of uncatalyzed reactions
a
decreases in magnitude with increase in temperature (-1.47,
-1.33, -1.28, -1.10 and -1.02 at 303, 308, 313, 318 and
323 K, respectively). The reaction constant, or sensitivity
constant, q?, describes the susceptibility of the reaction to
substituents and is a measure of the extent of electron demand
at the seat of reaction. The negative q? value in the present
study indicates that positive charge builds up in the transition
state. Asthevalueofq? isrelatedtothechargeduringtherate-
determining step (kd), it decreases with increase in
temperature.
The same stoichiometry was observed with all the
studied 4-oxo acids. Benzoic acid, after separation by
HPLC (m.p. 121 °C), was estimated quantitatively
(92–96% yield) with a standard curve at 235 nm (kmax).
CO2 was identified by bubbling N2 gas through the reaction
mixture and passing the liberated gas through a U-shaped
tube containing a saturated Ba(OH)2 solution. Formation of
a white precipitate of BaCO3 indicated the liberation of
CO2. A test for the formation of Br- was carried out by
adding silver nitrate solution to the reaction mixture after
completion of the reaction. The pale yellow precipitate of
silver bromide confirmed the formation of bromide as one
of the products of the reaction. No molecular bromine was
observed in these experiments.
Stoichiometry and product analysis
Reaction mixtures containing different ratios of [bromate]
to [oxo acid] in the presence of the requisite amounts of
other reagents were equilibrated at 40 °C for 12 h. The
unreacted [bromate] was determined by iodometry. The
results (Table 4) clearly demonstrate the consumption of
7 moles of bromate for 3 moles of the oxo acid. This
stoichiometry, together with the detection of benzoic acid,
bromide and CO2 in the product analysis, indicates the
following stoichiometric equation:
On the basis of the observed kinetic orders, the rate law
can be expressed as:
ꢂd½bromateꢃ
x
y
¼ k½bromateꢃ½Hþꢃ½oxo acidꢃ ½Ru(III)ꢃ
ð2Þ
dt
where x and y are less than unity and this equation holds
good for both oxo acid and Ru(III) in the studied concen-
tration range.
COOH
O
C-CH2-CH2-COOH
_
+ 7 BrO3
_
+ 9 CO2 + 7 Br + 6 H2O
3
3
R
R
123

Transition Met Chem (2011) 36:707–719
713
Fig. 4 a Arrhenius plots of log
k against 1/T. b Plot of DH#
against DS#. c Exner’s plot of
(a)
log k323 K against log k313 K
Reaction conditions as in
Table 3
.
A
k
E
103 / T ( K–1
)
120
100
80
60
40
20
0
(b)
(c)
E
E
D
D
C
C
B
k
B
#
A
H
A
-200
-150
-100
-50
0
50
S # ( J K–1 mol –1
)
5 + log
k (313 K)
Δ
ensuring that the oxidation takes place purely through
bromate.
Discussion
Acetate ion catalysis in the bromate–bromide reaction
was reported by Cortes and Faria [17], and the experi-
mental results were interpreted in a six-term rate law, with
and without involvement of acetate ion. The mechanism
proposed considers that when acetate binds to bromate, it
facilitates its second protonation. However, in the present
study, no catalytic effect of acetate was observed when
mercury acetate was added to the reaction mixture over a
wide concentration range (0.001–0.01 mol dm-3).
Role of mercury(II) acetate
The presence of mercury(II) (in excess of bromate con-
centration) simply means that Br2 oxidation is completely
suppressed; Br2 otherwise would have been formed by the
interaction between bromate and bromide (the reduction
product of bromate) as:
BrOꢂ3 þ 5Brꢂ þ 6Hþ ꢀ 3Br2 þ 3H2O
ð3Þ
No free radicals were observed in these reactions, and the
reaction rates were independent of [Hg]. These consider-
ations confirm that Hg(II) acts only as a scavenger [11, 14]
for any bromide formed, forming HgBr? (K1 = 2.51 9 108
dm3 mol-1) and HgBr2 (K2 = 3.80 9 108 dm3 mol-1) and
Identification of reactive species
Potassium bromate is a strong electrolyte, and in aqueous
solution, it exists as BrO3- ion, which is a very weak base
123

714
Transition Met Chem (2011) 36:707–719
4
3.5
3
Table 4 Consumption of bromate in the oxidation of 4-oxo acid
(a)
[Oxo acid] 9
103 (mol dm-3
[Bromate] 9
103 (mol dm-3
)
[Bromate]:
[Oxo acid]
)
Initial
Final
Consumed
1.0
2.0
5.0
10
20
50
7.67
15.35
38.26
2.33
4.67
2.33
2.32
2.35
11.74
2.5
2
k
Solution conditions: [Ru(III)] = 3.3 9 10-5 mol dm-3, [H2SO4] =
1.00 mol dm-3, [Hg(OAc)2] = 5. 00 9 10-3 mol dm -3 and AcOH–
H2O = 1:1% (v/v), Temp. 313 K
1.5
1
A study on oxidation states of Ru has shown [19–22]
that Ru(III) undergoes the following equilibria in acid
medium.
3ꢂ
RuCl3 þ 3HCl ꢀ ½RuCl6ꢃ þ 3Hþ
ð4Þ
ð5Þ
0.5
-0.8
-0.3
0.2
0.7
1.2
3ꢂ
½RuCl6ꢃ þ H2O ꢀ ½RuCl5ðH2O)]2ꢂ þ Clꢂ
From ion-exchange studies, it is confirmed that RuCl3
forms [RuCl6]3- in hydrochloric acid [19], and aquation of
these species to form [RuCl5(H2O)]2- requires only a few
seconds [20–22]. Therefore, it is assumed that Ru(III) in
the form of [RuCl5(H2O)]2- is taking part in the reaction.
The negative effect of Cl- ion on the rate of reaction is also
in support of this assumption. The same species was also
assumed to be the reactive catalyst in Ru(III)-catalyzed
oxidations of various organic substrates by different oxi-
dants [8, 23–25] in acid medium. In acid solutions,
[RuCl5(H2O)]2- shows an absorption maximum (kmax) of
490 nm. Neither bromate nor the 4-oxo acid has kmax in
this region. A change in kmax and the absorption of Ru(III),
on the addition of oxo acid (in the absence of bromate),
indicate the formation of a complex between Ru(III) and
the oxo acid, which is likely to be an outer sphere complex.
There is no absorption at 690 nm, which is characteristic of
Ru(V) species. Formation of the complex was also impli-
cated by a non-zero intercept in the plot of 1/k versus
1/[oxo acid] (Fig. 1). Similar complex formation between
organic substrates and Ru(III) has been observed in the
literature [8, 23–25]. The uncatalyzed acid bromate oxi-
dation of 4-oxo acids was first order in both [oxo acid] and
[bromate], and 1/kun versus 1/[oxo acid] plots were linear
passing through origin. The apparent fractional order with
respect to [Ru(III)] and a definite intercept in the 1/k versus
1/[Ru(III)] plots (Fig. 3) are further evidence for complex
formation between the oxo acid and Ru(III), which sub-
sequently reacts with bromate. The 4-oxo acids are weak
acids with pKa of 5.77 in aqueous solution at 40 °C [26];
hence, in acid medium, they will exist in the undissociated
form. 4-Oxo acids enolize in acid medium.
4
3.5
3
(b)
2.5
2
k
A
B
1.5
1
C
D
0.5
0
E
-1.5
-1
-0.5
0
0.5
1
Fig. 5 a Hammett plots of log k against r at different temperatures.
b Hammett plots of log k against r1 at different temperatures.
Reaction conditions as in Table 3. A 323 K, B 318 K, C 313 K,
D 308 K, E 303 K
[17]. HBrO3 is a very strong acid [18] and its protonation to
?
H2BrO3 is not significant under our experimental condi-
tions. Hence, in the present experiments, BrO3- is assumed
to be the main reactive species of the oxidant and the
?
involvement of HBrO3 or H2BrO3 can be excluded.
123

Transition Met Chem (2011) 36:707–719
715
Scheme 1 Reaction
OH
C
O
C
K
e
mechanism for the oxidation of
4-oxo acids by bromate in the
presence of Ru(III)
CH-CH2-COOH
CH2-CH2-COOH
H+
(Step 1)
R
R
R
Keto-form
Enol-form
H
O
O
OH
(III)
Ru
CH
K B
OH
C
C
C
CH
CH COOH + Ru(III)
2
CH
2
R
(Step 2)
Complex C1
H
O
H
O
(III)
O
OH
O
O
O
OH
Ru
Ru
(III)
KT
C
C
Br
+
C
C
_
CH
CH
H
C
CH
R
2
O
O
2
R
Br
O
(Step 3)
_
O
Complex C1
Complex C2
H
O
O
OH
+
Ru
(III)
C
COOH
C
kd
CH
CH
+
O
O
2
H
R
+
OHC CH COOH + HOBr + Ru(III)
2
r. d. s
Br
O
_
R
(Step 4)
fast
_
+
OHC
CH COOH + HOBr
2
CH
HOOC
COOH + H
+ Br
2
( O )
fast
HOOC CH COOH
2
CO2
+ other products
Reaction mechanism
then reacts with the binary complex C1, giving a cyclic
bromate ester (ternary complex C2). The concerted cis 1,3-
-
cycloaddition of BrO3 to the carbon–carbon double bond
The observed stoichiometry and effects of various param-
eters on the kinetics of this reaction allow us to propose a
mechanism as in Scheme 1.
of the substrate has also been proposed in the bromate
oxidation of cinnamic acids [27] and in the pyridinium
fluorochromate oxidation of 4-oxo acids [28]. Ru(III) ene
p-complex formation enhances the reactivity of the C=C
of complex C1 to form the cyclic bromate ester. The
decomposition of the five-membered ring complex C2
involving C–C and O–Br bonds (in the rate-determining
step), results in benzoic acid and formyl acetic acid, and
the oxidation state of Ru(III) is unaltered during the course
of the reaction. This is consistent with the fact that no
Ru(V) (kmax = 690 nm) or Ru(I) species were observed by
uv–vis. during the reaction. The formyl acetic acid on
oxidation with the HOBr product converts into malonic
acid along with the formation of Br- ion, further reactions
giving CO2 and H2O as final products. The intermediates
formed during the oxidation of malonic acid were identi-
fied by HPLC and by various tests [11]. The proposed
mechanism is also in accordance with the observed stoi-
chiometry. This mechanism is an example for the neigh-
boring group participation with intramolecular catalysis of
Ru(III) in the oxidation of 4-oxo acids.
The rate of enolization (determined by the bromination
method) is greater than the rate of oxidation by a factor of
ca. 15–18; hence, the step involving enol formation will not
be the rate-determining step and the reaction may be
visualized as proceeding via the enol form of the oxo acid.
In our proposed mechanism, the active species of Ru(III)
forms an outer sphere complex with the enol form of the
oxo acid (complex C1). The alkene, –OH and –COOH
groups could all be involved in complex formation with the
transition metal. Although evidence for specific interac-
tions was not obtained, the spectroscopic evidence for
formation of such a complex was obtained from the UV–
visible spectra of the catalyst, compared with a mixture of
catalyst and oxo acid in acid medium. A bathochromic shift
from 490 to 496 nm, with hyperchromic shift in the molar
extinction coefficient of the Ru(III) solution in the presence
of oxo acid, was obtained. Formation of the complex was
also implicated by a non-zero intercept in the plots of 1/
k against 1/[oxo acid] and 1/k against 1/[catalyst]. Bromate
123

716
Transition Met Chem (2011) 36:707–719
1
Rate law and verification
1
1
¼
þ
ð17Þ
k
kdKT KBKe½RuðIIIÞꢃ ½Oxo acidꢃ kd
On the basis of the reaction in Scheme 1, the rate in terms
of decrease in concentration of bromate can be expressed
as Eq. 6;
According to Eq. 16, other conditions being constant, a
plot of 1/k versus 1/[H?] should be linear passing through
the origin and the same is observed (Fig. 3). Similarly,
according to Eq. 17, plots of 1/k versus 1/[oxo acid] (at
constant [Ru(III)]) and 1/k against 1/[Ru(III)] (at constant
[oxo acid]) should be linear with a definite intercept on the
1/k axis. These observations (Figs. 1, 2) support the
validity of the rate law (Eq. 14). The proposed mechanism
is further supported by the solvent influence on the reaction
rate. The intermediate complex C2 is less polar than the
reactants due to dispersal of charge, and hence, decreasing
polarity of the solvent is expected to stabilize complex C2
in preference to the reactants, thereby enhancing the
reaction rate. Such a solvent influence has indeed been
observed (Table 1). The positive values of the free energy
and enthalpy of activation indicate that the transition state
is highly solvated, while the negative entropy of activation
suggests the formation of a compact activated complex
with fewer degrees of freedom.
ꢂd½BrOꢂ3 ꢃ
¼ kd½Complex C2ꢃ ½Hþꢃ
ð6Þ
dt
On the basis of the equilibrium steps in Scheme 1, Eqs. 7
and 8 can be obtained as follows:
½Complex C1ꢃ ¼ KBKe½Oxo acidꢃ ½RuðIIIÞꢃ
ð7Þ
ð8Þ
½Complex C2ꢃ ¼ KT KBKe½Oxo acidꢃ ½RuðIIIÞꢃ ½BrOꢂ3 ꢃ
With the help of Eqs. 6–8, we can write Eq. 9 as
ꢂd½BrOꢂ3 ꢃ
¼ kdKT KBKe½Oxo acidꢃ ½RuðIIIÞꢃ ½BrOꢂ3 ꢃ ½Hþꢃ
dt
ð9Þ
Since, bromate is present in complexed and uncomplexed
forms, at any time the total concentration of bromate can be
given as
½Br(V)ꢃT ¼ ½BrOꢂ3 ꢃ þ ½Complex C2ꢃ
ð10Þ
From the intercept and slope data of Figs. 1 and 2, the
decomposition constant (kd), and enolization and com-
plexation constants (Ke and KBKT) have been evaluated
and the values are presented in Table 5.
½Br(V)ꢃT ¼ ½BrOꢂ3 ꢃ þ KT KBKe½Oxo acidꢃ ½RuðIIIÞꢃ ½BrOꢂ3 ꢃ
ð11Þ
½Br(V)ꢃT ¼ ½BrOꢂ3 ꢃ þ ½1 þ KT KBKe½Oxo acidꢃ ½RuðIIIÞꢃꢃ
Structural influences and linear free-energy
relationships
ð12Þ
Therefore,
The order of reactivity among the phenyl substituted
4-oxo acids is p-methoxy ꢁ p-methyl [ –H [ p-chloro [
m-nitro oxo acid. A plot of log k versus Hammett’s sigma
values gave a smooth curve at all the studied temperatures
(Fig. 5a). The curvature may be attributed to (1) a gradual
change in the reaction mechanism when one passes from
½BrðVÞꢃT
½BrOꢂ3 ꢃ ¼
ð13Þ
1 þ KT KBKe½Oxo acidꢃ ½RuðIIIÞꢃ
The rate law in terms of total bromate concentration can
then be given as
ꢂd½BrðVÞꢃT kdKT KBKe½Hþꢃ ½Oxo acidꢃ ½RuðIIIÞꢃ ½BrðVÞꢃT
¼
dt
1 þ KT KBKe½RuðIIIÞꢃ ½Oxo acidꢃ
Table 5 Evaluated decomposition constant (kd), enolization and
formation constants (KeKBKT) of Ru(III)-catalyzed oxidation of 4-oxo
acids at 313 K using Eq. 17
ð14Þ
The rate law (Eq. 14) is in accordance with the experi-
mental results, wherein a first-order dependence on both
[bromate] and [acid], and fractional orders in both [oxo
acid] and [catalyst] were observed.
Substituent
kd 9 103 s-1
(decomposition constant)
KeKBKT 9 10-5
(formation constants)
p-OCH3
p-CH3
–H
22.72
6.90
3.12
1.64
0.45
37.42
25.69
22.28
21.37
21.06
The rate law (Eq. 14) may be rearranged to Eqs. 15–17,
which are suitable for verification.
Rate
kdKT Ke½Hþꢃ ½RuðIIIÞꢃ ½Oxo acidꢃ
p-Cl
¼ k ¼
ð15Þ
½BrðVÞꢃT
1 þ KT KBKe½RuðIIIÞꢃ ½Oxo acidꢃ
m-NO2
ꢀ
ꢁ
Solution conditions: [Oxo acid] = 1.00 9 10-2 mol dm-3, [bro-
1
1
1
1
¼
þ
ð16Þ
mate] = 1.00 9 10-3 mol dm-3, [Ru(III)] = 3.3 9 10-5 mol dm-3
,
k
kdKT KBKe½RuðIIIÞꢃ ½Oxo acidꢃ kd ½Hþꢃ
[H2SO4] = 1.00 mol dm-3, [Hg (OAc)2] = 5. 00 9 10-3 mol dm-3
and AcOH–H2O = 1:1% (v/v)
At constant acid (1.0 mol dm-3) concentration, Eq. 16
reduces to Eq. 17
The average error in kd and KeKBKT values is 5%
123

Transition Met Chem (2011) 36:707–719
717
electron-donating to electron-withdrawing substituents, (2)
a change in the rate-determining step or (3) a change in the
nature of transition state. The linearity of the Exner’s plot
(Fig. 4c) suggests that these reactions follow a similar
mechanism [29]. The isokinetic temperature (b) was cal-
culated as 370.2 K from the slope of the plot of log k at
323 K against log k at 313 K (Fig. 4c), indicating that the
rate is governed by the enthalpy of activation. This is also
corroborated by near constancy in DG# values.
decreases the rate of oxidation of meta-substituted com-
pounds through the inductive effect. But the mesomeric
effect in para-substituted compounds overweighs the
inductive effect, resulting in overall electron release to the
aromatic ring. Thus, in the oxidation of p-methoxy 4-oxo-
4-phenylbutanoic acid, direct mesomeric interaction with
the reaction center also comes into play, resulting in greater
electron release than expected on the basis of the Hammett
r value.
From the preceding discussion, it may be construed that
the curve in plot of log k versus r (Fig. 5a) is a result of the
change in the nature of the transition state when one goes
from electron-releasing to electron-withdrawing substitu-
ents, i.e., an electronic perturbation effect. The transition
state (Scheme 1) involves both C–C bond cleavage and C–
O bond formation. If both the steps are influenced in the
same way by the substituents, a linear Hammett plot is
expected. However, if the two processes influence the
reaction in different ways, a curved Hammett plot will be
obtained. Electron-releasing substituents may lead to a
transition state, in which C–C bond-breaking has pro-
gressed to a greater extent than the C–O bond formation,
while the converse would be true for electron-attracting
substituents. The activation parameters are in favor of this
view. The DH# value is seen to decrease with increasing
electron-donating ability of the substituents (Table 3),
whereas, DS# becomes more negative, reflecting a more
solvated transition state. For electron attractors like the
nitro group, bond formation is apparently more important,
since DH# increases, resulting in decreased rate of solvol-
ysis, while DS# becomes positive. Higher values of sigma,
Isokinetic relationships and catalytic efficiency
The Hammett reaction constant (q?) decreases in magni-
tude with increase in temperature, from -1.47 at 303 K to
-1.02 at 323 K. This indicates that the carbocationic
character of the transition state decreases with increase in
temperature. At the isokinetic temperature, q? becomes
zero, and there is no sensitivity to substituents. The Ham-
mett lines intersect at a point corresponding to an r1iso value
of -1.20 and a kiso value of 31.62 9 10-3 s-1 (Fig. 5b).
A reaction series that exhibits a common point of
intersection in the Hammett plot (Fig. 5b), is expected to
have a common point of intersection in the Arrhenius plot
also [29]. The Arrhenius plots are shown in Fig. 4a. From
the plots of DH# against DS# (Fig. 4b) and q? against 1/T
(Fig. 6), according to the equation
q ¼ qað1 ꢂ b=TÞ
ð18Þ
-0.9
-1
?
i.e., exalted Brown and Okamoto’s r values [16] can be
used to give linear correlations. This is observed in the
present study; a plot of log k versus r? was linear with a
slope of -1.28 at 313 K (Fig. 5b). The reaction constant
(q?), which is a measure of the extent of electron demand
at the seat of reaction, decreases in magnitude with
increase in temperature. The negative value of q? suggests
the sensitivity of the reaction to the effects of electronic
perturbation. It also provides information about the nature
of the transition state. A reaction involving positive charge
development in the transition state is aided by electron-
releasing substituents and the q? value will be negative
[30], as in the present case. The q? value (-1.02 to -1.47)
in the present study may be attributed to the nature of the
observed rate constant (k), which is a composite of several
terms like enolization, complexation, oxidation, etc.
A
-1.1
B
ρ+ -1.2
-1.3
C
D
-1.4
E
The effect of various substituents on the reaction rate is
as expected, with the exception of p-methoxy, which
showed a significant increase in the rate. This may be
explained on the basis of the r values of the methoxy
substituent. The rm value for methoxy is ?0.115, while rp
is -0.268. These values show that the methoxy group
-1.5
3.05
3.1
3.15
3.2
3.25
3.3
3.35
103 / T ( K-1 )
Fig. 6 Plot between q? and 1/T at different temperatures. Reaction
conditions as in Table 3. A 323 K, B 318 K, C 313 K, D 308 K,
E 303 K
123

718
Transition Met Chem (2011) 36:707–719
Table 6 Values of catalytic constant (KC) at different temperatures and activation parameters calculated using KC value
4-oxo acid
KC 9 102
Ea
(kJ mol-1
DH#
DS#
DG#
)
(kJ mol-1
)
(J mol-1 K-1
)
(kJ mol-1
)
298
303
12.77
308
16.71
313
22.28
318
33.43
323 K
–H
8.35
111.11
20.06
3.33
58.92
(–)
53.16
29.45
58.50
80.84
107.51
51.01
26.85
55.90
78.24
104.91
-94.9
-153.8
-72.7
-11.8
60.4
80.72
75.01
78.66
60.4
p-OCH3
p-CH3
p-Cl
133.70
27.80
5.55
167.1
39.12
7.59
200.5
50.15
13.91
3.00
222.8
78.02
23.40
6.12
111.45
38.98
11.11
m-NO2
0.51
1.00
1.67
85.98
Solution conditions: [Oxo acid] = 1.00 9 10-2 mol dm-3
,
[bromate] = 1.00 9 10-3 mol dm-3
,
mol dm-3 [Ru(III)] = 3.3 9 10-5
,
mol dm-3, [H2SO4] = 1.00 mol dm-3, [Hg (OAc)2] = 5. 00 9 10-3 mol dm-3 and AcOH–H2O = 1:1% (v/v)
The average error in values of Ea, DH#, DS# and DG# are 1.6 kJ mol-1
,
1.6 kJ mol-1
,
3 J mol-1 K-1 and 2 kJ mol-1, respectively. DH#,
DS# and DG# values were obtained at 313 K. The average error in the catalytic constant is 5%
Conclusion
The Tiso value is 369.0 K, in good agreement with the
value of 370.2 K obtained from the Exner’ s plot (Fig. 4c).
From the Hammett and Arrhenius plots associated with
Moelwyn-Hughes [31] pointed out that when catalyzed
isokinetic relationships, the q? values are negative and
decrease with increasing temperature. Electron donors
accelerate the reaction rate and electron acceptors retard it.
The structure-reactivity correlations suggest a carbocat-
ionic reaction center in the transition state. It may be
concluded that Ru(III) acts as an efficient catalyst in the
oxidation of 4-oxo acids by bromate in acid medium.
Unlike earlier studies, this reaction is an example of the
neighboring group participation and intramolecular catal-
ysis and is a valuable protocol for the synthesis of substi-
tuted benzoic acids.
and uncatalyzed reactions proceed simultaneously, the
relationship is given as:
y
kobs ¼ kun þ KC½catalystꢃ
ð19Þ
where kobs is the observed pseudo-first-order rate constant
obtained in the presence of Ru(III) catalyst, kun is that for
the uncatalyzed reaction, KC is the catalytic constant and y
is the order of the reaction with respect to [Ru(III)], which
is 0.5 in the present study. The value of KC was calculated
using the equation
0:5
KC ¼ ðkobs ꢂ kunÞ=½RuðIIIÞꢃ
ð20Þ
Acknowledgments We wish to thank the anonymous reviewer and
editor for their critical and useful comments which refined the man-
uscript a lot. Financial assistance from the University Grants Com-
mission, New Delhi, India, under the Special Assistance Programme
(No. F. 540/6/DRS/2009/SAP-I) is gratefully acknowledged.
The values of KC were calculated at different temperatures
(298–323 K), and a plot of log KC versus 1/T was linear
(R2 C 0.993). The energy of activation and other thermo-
dynamic parameters for the catalyst are tabulated in
Table 6. The high KC value indicates that Ru(III) is an
efficient catalyst in the bromate oxidation of 4-oxo acids.
Finally, we have compared our findings with those
reported for the osmium(VIII) catalyzed oxidation of oxo
acids by alkaline hexacyanoferrate(III). The proposed
mechanism in this case involves rapid complex formation
between the enolate anion of the 4-oxo acid and [OsO4-
(OH)2]2-, which decomposes into products in the slow step
giving Os(VI). The Os(VI) is reoxidized to Os(VIII) by
[Fe(CN)6]3- in a fast step [32]. On the other hand, in the
alkaline hexacyanoferrate(III) oxidation of 4-oxo acids, a
mechanism involving the formation of enolate anion from
the oxo compound and subsequent rate-determining elec-
tron transfer by the oxidant has been proposed [33]. In the
present investigation, the experimental results are entirely
different from the earlier reports and a different mechanism
wherein neighboring group participation in the intramo-
lecular catalysis is proposed.
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