The kinetics and the mechanism of oxidative decarboxylation of benzilic acid by acidic permanganate (stopped-flow technique) - An autocatalytic study

Article abstract of DOI:10.1139/v04-102

The kinetics of the oxidation of benzilic acid by potassium permanganate in an acidic medium were studied spectrophotometrically. The reaction followed a two-stage process, wherein both stages of the reaction followed first-order kinetics with respect to permanganate ion and benzilic acid. The rate of the reaction increased with an increase in acid concentration. Autocatalysis was observed by one of the products, i.e., manganese(II). A composite mechanism involving autocatalysis has been proposed. The activation parameters of the reaction were calculated and discussed and the reaction constants involved in the mechanisms were calculated. There is a good agreement between the observed and calculated rate constants under different experimental conditions.

Full text of DOI:10.1139/v04-102

1372
The kinetics and the mechanism of oxidative
decarboxylation of benzilic acid by acidic
permanganate (stopped-flow technique) — An
autocatalytic study
Sairabanu A. Farokhi and Sharanappa T. Nandibewoor
Abstract: The kinetics of the oxidation of benzilic acid by potassium permanganate in an acidic medium were studied
spectrophotometrically. The reaction followed a two-stage process, wherein both stages of the reaction followed first-
order kinetics with respect to permanganate ion and benzilic acid. The rate of the reaction increased with an increase
in acid concentration. Autocatalysis was observed by one of the products, i.e., manganese(II). A composite mechanism
involving autocatalysis has been proposed. The activation parameters of the reaction were calculated and discussed and
the reaction constants involved in the mechanisms were calculated. There is a good agreement between the observed
and calculated rate constants under different experimental conditions.
Key words: oxidation, autocatalysis, benzilic acid, two-stage kinetics.
Résumé : Faisant appel à la spectrophotométrie, on a étudié la cinétique de l’oxydation de l’acide benzilique par le
permanganate de potassium en milieu acide. La réaction se produit en deux étapes et chacune d’elles est du premier
ordre par rapport aux concentrations de l’ion permanganate et de l’acide benzilique. La vitesse de la réaction augmente
avec une élévation de la concentration d’acide. On a observé de l’autocatalyse par l’un des produits, soit le manga-
nèse(II). On propose un mécanisme composé impliquant de l’autocatalyse. On a calculé et on discute des paramètres
d’activation de la réaction; on a aussi calculé les constantes de réaction impliquées dans les mécanismes. On a observé
un bon accord entre les constantes de vitesses observées et calculées dans différentes conditions expérimentales.
Mots clés : oxydation, autocatalyse, acide benzilique; cinétique à deux étapes.
[Traduit par la Rédaction] Farokhi and Nandibewoor 1380
Introduction
mediate complex formation, and in others, the results are in-
terpreted by a free radical mechanism. The mechanisms sug-
gested by various authors are not uniform, indicating that a
wide variety of mechanisms are possible depending on the
nature of the reactive species of permanganate as well as the
substrate.
Several studies have been reported on the oxidation of
benzilic acid by other oxidants such as cerium(IV) (12, 13),
ferricyanide (14), chromic acid (15), thallium(III) (16), per-
sulphate (17), and N-bromosacharin (18). Even though the
oxidation product of benzilic acid is the same in both acidic
and alkaline media, the reactions differ in many kinetic pa-
rameters. Interest in kinetic studies, in particular of α-
hydroxy acids (19), has developed and no attempts have
been made to correlate the rate data with the structural dif-
ferences in these acids.
It is evident that manganese(VII) is reduced to various ox-
idation states in acidic, alkaline, and neutral media during
oxidation by permanganate. Out of the six oxidation states
of manganese (1), from 2+ to 7+, permanganate Mn(VII), is
the most powerful oxidant in a dilute acid medium with re-
duction potentials of 1.69 V for the Mn(VII)/Mn(IV) couple
and 1.51 V for the Mn(VII)/Mn(II) couple.
Oxidations by permanganate ion are applied extensively in
organic syntheses (2–8), especially since the advent of
phase-transfer catalysis (4, 5, 7). Kinetic studies are impor-
tant sources of mechanistic information on the reactions, as
demonstrated by the results for unsaturated acids both in
aqueous (2, 4, 8) and nonaqueous media (9).
Furthermore, the mechanism by which this multivalent
oxidant oxidizes a substrate depends not only on the sub-
strate but also on the medium (10, 11) used for the study. In
some cases the mechanistic approach is based on the inter-
In view of the lack of reports in the literature on the oxi-
dation of benzilic acid by acidic KMnO4, we have selected it
as a substrate for study. The present study deals with the ki-
Received 20 February 2004. Published on the NRC Research Press Web site at http://canjchem.nrc.ca on 23 October 2004.
S.A. Farokhi and S.T. Nandibewoor.1 P.G. Department of Studies in Chemistry, Karmatak University, Dharwad, 580 003, India.
1Corresponding author (e-mail: stnandibewoor@yahoo.com).
Can. J. Chem. 82: 1372–1380 (2004)
doi: 10.1139/V04-102
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Farokhi and Nandibewoor
1373
netics of the oxidation of benzilic acid by KMnO4 in an
acidic medium to arrive at a plausible mechanism and to de-
termine the active species of permanganate in such media.
60% to 80% completion of the reaction. kobs values were
reproducible within 5%, where the first slow stage is fol-
lowed by a second fast stage process (Fig. 1). This type of
two-stage process has been observed in a few cases (23, 24).
During the progress of the reaction, the colour of the solu-
tion changed from violet to colourless, indicating the forma-
tion of Mn(II), which was verified by a spot test (20 b). The
formation of Mn(II) is shown spectrally in Fig. 2.
Experimental
Since the initial reaction was too fast to monitor by the
usual method, kinetic measurements were performed on a
Hitachi 150-20 spectrophotometer connected to a rapid ki-
netic accessory (HI-TECH SFA-12).
Results
Materials and methods
Stoichiometry and product analysis
All chemicals used were of BDH Analar specifications.
Potassium permanganate (BDH) was prepared in doubly dis-
tilled water and its strength was ascertained (20a) by titrat-
ing against standard H2(CO2)2 solution in hot acidic
medium. The solubility of benzilic acid (Merck) (BA) is low
and it was prepared in water by warming. Perchloric acid
(Merck) and sodium perchlorate (BDH) were used to pro-
vide the required acidity and ionic strength, respectively.
Mn(II) was prepared by dissolving manganese(II) sulphate
(Riedel). Doubly distilled conductivity water was used to
prepare all the solutions.
The optical density vs. concentration plot for acidified
permanganate ion showed that Beer’s law, under the reaction
conditions in 6.0 × 10–5 to 2.0 × 10–4 mol dm–3 in
0.01 mol dm–3 HClO4, was obeyed at a wavelength of
526 nm and the molar absorptivity of permanganate was
found to be 2300 50 dm3 mol–1 cm–1, which was in good
agreement with previous results (21, 22).
Different sets of reaction mixtures containing excess
–
[MnO4 ] over [benzilic acid] were mixed in the presence of
1.0 × 10–3 mol dm–3 perchloric acid, adjusted to an ionic
strength of 5.0 × 10–3 mol dm–3, and kept for 12 h in an inert
atmosphere at 21 °C. After completion of the reaction, the
remaining permanganate was then determined spectrophoto-
metrically. The results indicated that 5 mol of benzilic acid
were consumed by 2 mol of permanganate as in eq. [1].
The main oxidation products were identified as benzophe-
none, CO2, and manganese(II). The product, manganese(II)
in acid medium, was analyzed using EDTA solution with
Erichrome black-T as the indicator (20b). The formation of
benzophenone was confirmed by its IR and NMR spectra
(25). Further, the yield was quantitatively estimated to ca.
85%, evidence for which is provided by isolation of its 2,4-
DNP derivative (26), with a mp of 236 °C. Benzophenone
did not undergo further oxidation under the stated kinetic
conditions. The test for the probable oxidation product of the
ketone, i.e., benzoic acid, was negative.
Kinetic measurements
All kinetic measurements were performed under pseudo-
first-order conditions where [benzilic acid] was always in
Reaction orders
The reaction order for benzilic acid and perchloric acid
was investigated by varying one of these concentrations
while keeping all other concentrations and conditions con-
stant and studying the effect of its concentration on the rate
of the reaction.
–
excess over [MnO4 ] at a constant ionic strength of 5.0 ×
10–3 mol dm–3 in acidic medium at a constant temperature of
21 0.1 °C.
The progress of the reaction was followed by monitoring
–
–
the decrease in MnO4 in a 1 cm quartz cell of a Hitachi
Kinetic data were collected for initial [MnO4 ] in the
model 150-20 spectrophotometer at its absorption maximum
(526 nm) as a function of time. Earlier, it was verified that
there is negligible interference from other reaction species at
this wavelength. The pseudo-first-order rate constants were
obtained from the plots of log (At – A∞) vs. time and the plot
of log [(At – A∞) – (At′ – A∞)] vs. time, where At and A∞ rep-
resent the absorbance at time t and ∞, respectively. At′ repre-
sents the absorbance for the extrapolated point at time t′. The
plots of log (At – A∞) vs. time and of log [(At – A∞) – (At′ –
A∞)] vs. time indicate that there was a two-stage process and
that both stages were linear. In the first slow process, the lin-
earity was observed up to 60% completion of the reaction
and in the second fast process, linearity was observed from
(0.40–4.0) × 10–4 mol dm–3 range. In an acid medium, the
reaction followed a two-stage process (Fig. 1). All the exper-
iments exhibited an identical pattern, with the initial slow re-
action followed by a faster one. The plots of log (At – A∞)
vs. time and log [(At – A∞) – (At′ – A∞)] vs. time (r > 0.9355,
s ≤ 0.0144 for stage 1, r > 0.9951, s ≤ 0.4480 for stage 2) for
–
different initial concentrations of MnO4 exhibit two distinct
straight lines for the slow and fast stages of the reaction
(Fig. 1). The fairly constant ks (s–1) values for the slow stage
and kf (s–1) values for the fast stage confirms that slow and
–
fast depletions in MnO4 follow first-order kinetics. The
concentration of benzilic acid was varied in the range of
(0.40–4.0) × 10–3 mol dm–3, keeping all other reactant con-
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Can. J. Chem. Vol. 82, 2004
–
Fig. 1. Plot of (a) log(At – A∞) vs. time and (b) log[(At – A∞) – (At′ – A∞)] vs. time in an acidic medium ([MnO4 ] = 2 ×
10–4 mol dm–3, [BA] = 2 × 10–3 mol dm–3, [H+] = 1.0 ×10–3 mol dm–3, and I = 5.0 × 10–3 mol dm–3.
[H+] on the rate of the reaction, the [H+] was varied from
Fig. 2. Spectroscopic changes occurring in the oxidation of
(0.50–5.0) × 10–2 mol dm–3 at a constant concentration of
–
benzilic acid by permanganate with [MnO4 ] = 2.0 ×
–
10–4 mol dm–3, [BA] = 2.0 × 10–3 mol dm–3, [H+] = 1.0 ×
10–3 mol dm–3, and I = 5.0 × 10–3 mol dm–3 at 21 °C (scanning
time interval = 1 min).
MnO4 and benzilic acid, maintaining a constant ionic
strength of 5.0 × 10–3 mol dm–3. It was found that in both
stages the rate increased with an increase of [H+]. The order
in both stages was less than unity (Table 1). It was found
that the ionic strength and dielectric constant of the medium
have no significant effect on the rate of reaction.
Effect of added products
The addition of the product benzophenone did not show
any significant effect on the rate of the reaction. However,
the initial addition of the product Mn(II) in the range of
(0.20–2.0) × 10–4 mol dm–3, keeping other conditions con-
stant, had an appreciable increasing effect on the rate of the
reaction. As the initial concentration of Mn(II) was in-
creased, the rate progressively increased at an order of
around 0.62 and 0.26 in the slow and fast stages, respec-
tively. This illustrates the autocatalytic nature of the product
(Table 2), which is also evident from the concentration vs.
time plot (Fig. 3) and the linear plot of kobs vs. kcal for man-
ganese(II) variation (Fig. 3, inset), as observed in earlier
cases (27).
Polymerization study
centrations and conditions constant. In an acidic medium,
the ks and kf values (Table 1) increased with increasing con-
centration of benzilic acid indicating first-order dependence
on the [substrate] in both the stages. To study the effect of
To test for free radical intervention, the reaction mixture
–
containing [MnO4 ] = 2 × 10–4 mol dm–3 and [BA] = 2 ×
10–3 mol dm–3 were taken, and the pH of the solution was
adjusted to 3.0 The solution was then degassed with nitrogen
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Farokhi and Nandibewoor
1375
Table 1. Effect of [benzilic acid], [MnO4 ], and [H+] on the oxidation of benzilic acid by acidified permanganate at 21 °C, I = 5.0 ×
–
10–3 mol dm–3.
[H+]
ks
ks
kf
kf
–
[BA]
[MnO4 ]
(×103 mol dm–3) (×104 mol dm–3) (×103 mol dm–3) (×103 s–1) (Expt.) (×103 s–1) (Calcd.) (×103 s–1) (Expt.) (×103 s–1) (Calcd.)
5.21
13.2
5.01
12.5
0.4
1.0
2.0
3.0
4.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
0.4
1.0
2.0
3.0
4.0
2.0
2.0
2.0
2.0
2.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
0.5
0.8
1.0
2.0
5.0
0.41
1.12
2.07
3.10
4.23
2.08
2.03
2.07
2.16
2.12
1.15
1.81
2.07
3.60
5.85
0.42
1.05
2.10
3.16
4.22
2.10
2.10
2.10
2.10
2.10
1.16
1.75
2.10
3.57
6.11
27.5
37.2
51.3
27.3
27.5
27.5
27.9
28.0
14.3
21.0
27.5
36.4
60.2
25.1
37.6
50.1
25.1
25.1
25.1
25.1
25.1
14.5
21.2
25.1
39.0
60.2
Table 2. Effect of product Mn(II) on the oxidation of benzilic
dant undergoes five equivalent changes. In this experiment,
–
acid by acidified KMnO4 at 294 K ([BA] = 2.0 × 10–3 mol dm–3,
the reaction between benzilic acid and MnO4 had a stoichi-
–
[MnO4 ] = 2.0 × 10–4 mol dm–3, [H+] = 1.0 ×10–3 mol dm–3, and
ometry of 5:2 with first-order dependence on [MnO4 ] and
–
I = 5.0 × 10–3 mol dm–3).
[BA] and fractional order dependence on both [H+] and
[Mn(II)]. Stoichiometric studies revealed 5-equivalent
changes in Mn(VII). Hence, the formation of Mn(IV) as a
reduced product of Mn(VII) in the reaction was excluded,
although formation of Mn(IV) had been observed in earlier
studies (29, 30). The strong reducing character of benzilic
acid was evident in the reduction of Mn(VII) to Mn(II) in
the HClO4 medium.
[Mn(II)]
ks
kf
(×104 mol dm–3)
(×103 s–1)
(×103 s–1)
2.07
2.75
6.13
8.11
10.2
11.3
0.0
0.2
0.6
1.0
1.5
2.0
27.5
28.4
34.0
43.4
48.3
49.1
The active species of permanganate in an aqueous acid
medium may be deduced from the dependence of the rate on
[H+] in the reaction medium. The apparent order of less than
unity in [H+] may be an indication of the permanganate spe-
cies as permanganic acid. In an acidic medium, permanganic
acid (HMnO4) is a more efficient oxidant species of manga-
nese(VII) than permanganate ion (29–31). In addition, it was
observed that the reaction rate increased with an increase in
[H+] and tended to attain a limiting value at high acidities.
The plot of kobs vs. [H+] in an acidic medium is a curve of
decreasing slope (convex to the rate axis) and passing
through the origin (Fig. 4) from which it can be inferred (32)
that a rapid equilibrium between the protonated and unpro-
tonated form is involved. At higher acidities protonation is
almost complete, leading to the limiting rate, which indi-
cates that only the protonated form is active. The negligible
ionic strength effect on the rate of the reaction also confirms
and acrylonitrile (5 mL, 20% v/v) was added to the solution
at room temperature and was kept for 24 h. On diluting the
reaction mixture with methanol, no precipitate resulted indi-
cating no polymerization of acrylonitrile. This observation
indicates that the oxidation of benzilic acid by acidic potas-
sium permanganate proceeds through the absence of poly-
mer formation and, as well, any intervention of free radicals.
There is evidence of polymerization with acrylonitrile in the
earlier studies (28).
Effect of temperature
The effect of temperature on the reaction was studied at
four different temperatures. The rate increased with an in-
crease in temperature. The ∆H‡ and ∆S‡ values were
obtained from a log(kobs/T) vs. 1/T plot (r > 0.9719, s ≤
0.01298 for stage 1, r > 0.9839, s ≤ 0.07748 for stage 2).
The activation parameters ∆H‡ (kJ mol–1), ∆S‡ (J K–1 mol–1),
and ∆G‡ (kJ mol–1) were found to be 64 2, –79 6, and
87 4 for stage 1 and 28 2, –181 16, and 81 4 for
stage 2, respectively.
–
HMnO4 as the active species of MnO4 . Thus, the acid per-
manganate equilibrium is represented as
K
1
–
[2]
H+ + MnO4
HMnO4
ꢀ
where K1 is the formation constant of HMnO4.
Based on the experimental results, the following mecha-
nism can be proposed in which all the observed orders with
respect to each constituent such as [oxidant], [reductant],
and [H+] may be well accommodated. The results suggest
the formation of a protonated form of permanganate as the
Discussion
Oxidation of benzilic acid by permanganate in a HClO4
medium is a noncomplementary reaction in which the oxi-
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Can. J. Chem. Vol. 82, 2004
Fig. 3. Concentration vs. time plots to show autocatalysis. (a) [BA] = 2.0 × 10–3 mol dm–3, [MnO4 ] = 2.0 × 10–4 mol dm–3, [H+] =
–
1.0 × 10–3 mol dm–3, I = 5.0 × 10–3 mol dm–3 (b) [BA] = 3.0 × 10–3 mol dm–3, [MnO4 ] = 2.0 × 10–4 mol dm–3, [H+] = 1.0 ×
–
– 1
10–3 mol dm–3, I = 5.0 × 10–3 mol dm–3. Inset: plot of kexpt vs k′cal = [BA]1[Mn(II)]0.62[H+]0.70[MnO4 ] for stage 1 and k′′cal
=
– 1
[BA]1[Mn(II)]0.26[H+]0.6[MnO4 ] for stage 2 for the variation of manganese(II) (Table 2).
–
active species of MnO4 in a prior equilibrium, which reacts
with 1 mol of benzilic acid in a rate-determining step to give
a complex (C1) that further decomposes in a fast step to give
benzophenone and the intermediate Mn(V). The intermedi-
ate Mn(V), being unstable in an acidic medium, reacts with
another mol of benzilic acid in a fast step to Mn(III), which
is subsequently reduced to the end product Mn(II) in further
fast steps satisfying stoichiometric observations. The spec-
tral evidence for the formation of complex C1 was obtained
from UV–vis spectra of both benzilic acid and the mixture
of benzilic acid and permanganic acid, which indicated a
bathochromic shift from 325 to 347 nm. The steps proposed
are shown in Scheme 1.
Since none of the intermediates of Mn(VII) could be de-
tected, Scheme 1 is only one of the possible mechanisms for
the reaction in the absence of free radicals. Attempts were
made to allow spectrophotometric detection (in terms of
wavelengths) of Mn5+ and Mn3+ intermediates as the reac-
tion proceeded in the oxidation of benzilic acid by HMnO4.
Unfortunately, the low concentration of Mn3+ and Mn5+ in-
termediates obtained under our experimental conditions
made the detection unsuccessful. However, there are reports
of their existence in the earlier literature (33, 34). Scheme 1
leads to the rate law (eq. [3] that explains all the observed
orders except autocatalysis.
d[Mn(VII)]
k1K1[H+ ]T[Mn(VII)]T[BA]
[3]
[4]
−
=
=
dt
(1 + K1[H+ ])(1 + K1[Mn(VII)])
k1K1[H+ ][BA][Mn(VII)]
1 + K1[H+ ] + K1[Mn(VII)]) + K12[H+ ][Mn(VII)]
The terms such as K1[Mn(VII)] and K12[H+][Mn(VII)] in the
denominator of eq. [4] are negligibly small compared with
unity, in view of the low concentration of [Mn(VII)] used in
the experiments. Hence, eq. [4] becomes eq. [5]
Rate
k1K1[H+ ][BA]
1 + K1[H+ ]
[5]
kobs
=
=
[Mn(VII)]
The mechanism of Scheme 1 and rate law (eq. [3]) may be
verified by rearranging to the form
© 2004 NRC Canada

Farokhi and Nandibewoor
1377
Fig. 4. The plot of kobs vs. [H+] (a curve of decreasing slope (convex to the rate axis)).
Scheme 1.
[BA]
1
1
of this plot lead to values of K1 and k1 at 21 °C of 221
[6]
=
+
9 dm3 mol–1, 5.8 0.3 dm3 mol–1 s–1 for stage 1 and 371
13 dm3 mol–1, 46 2 dm3 mol–1 s–1 for the fast stage 2, re-
spectively. The experimental rate constants can be computed
for both stages using these values. There is a reasonable
agreement between the calculated and experimental rate
constants (Table 1). The negligible effect of ionic strength
kobs
k1K1[H+ ] k1
where kobs is first-order rate constant. According to eq. [6],
the plot of [BA]/kobs vs. 1/[H+] (r > 0.999, s ≤ 0.06990 for
stage 1, r > 0.9946, s ≤ 0.06993 for stage 2) is expected to
be linear, which is verified in Fig. 5. The slope and intercept
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1378
Can. J. Chem. Vol. 82, 2004
Fig. 5. The plot of [BA]/kobs vs. 1/[H+] (verification of rate law (eq. [3])).
Scheme 2.
present case, Mn(II) was found to catalyse the rate of oxida-
tion of benzilic acid (Table 2).
and relative permitivity on rate is qualitatively consistent
with the reaction between two neutral molecules as in
Scheme 1.
The apparent order of less than unity with respect to the
Mn(II) ion may be attributed to a complex formation be-
tween benzilic acid and Mn(II). The complex (C2) is then
subsequently involved in the interaction with HMnO4. These
steps shown in Scheme 2 will form a part of Scheme 1.
The evidence for complex formation (C2) was obtained
from UV–vis spectra of both benzilic acid and mixtures of
Autocatalysis by Mn(II) in an acidic medium
The autocatalysis by one of the products (Mn(II)) is inter-
esting. The effect of Mn(II) as a catalyst in the case of the
oxidation of permanganate is well-known (35–37). In the
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Farokhi and Nandibewoor
1379
benzilic acid and Mn(II), which indicated a hypsochromic
shift of about 4 nm from 338 to 334 nm and a new band ap-
peared at 372 nm. Indeed, such complex formation between
substrates such as styrene and Mn(II) has been observed in
the literature (38, 39).
media are consistent with product, mechanistic, and kinetic
studies.
Acknowledgement
Thus, when Mn(II) is initially present, a composite
scheme involving all the steps of Schemes 1 and 2 operates
and the rate law becomes
One of the authors, Sairabanu A. Farokhi, would like to
acknowledge the University Grants Commission (UGC),
Southwestern Regional Office, Bangalore, for providing a
Fellowship.
kgross = kobs + kautocat
k1K1[BA][H+ ]
kgross
=
1 + K1[H+ ]
References
[7]
k2K2[Mn(II)][BA]2
{1 + K2[BA]2}{1 + K2[Mn(II)][BA]}2
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+
kautocat = kgross − kobs
k2K2[Mn(II)][BA]2
{1 + K2[BA]2}{1 + K2[Mn(II)][BA]}2
[8]
kautocat
=
where kgross and kautocat refer to the rate paths of the overall
and autocatalytic paths, respectively. At constant concentra-
tions of oxidant and substrate, a plot of [BA]2/kautocat in
eq. [8] vs. 1/[Mn(II)] (r > 0.9987, s ≤ 0.1566 for stage 1, r >
0.9596, s ≤ 0.1574 for stage 2) for both stages was found to
be linear. Indeed, it is to be noted that the plot shows an in-
tercept that is in agreement with the complex formation as
shown in Scheme 2. The high negative values of ∆S‡ indi-
cate that complex (C1) is more ordered than the reactants
(40).
Reason for the two stages
Details regarding the occurrence of two successive stages,
slow and fast, might be due to the creation of an optimum
concentration of some intermediates (41–43). It might be
12. V.K. Grover, Y.K. Gupta, and K.J. Yugul. J. Inorg. Nucl.
Chem. 31, 1403 (1969).
13. S.B. Hanna and S.A. Sarac. J. Org. Chem. 42, 2063 (1977).
14. N.C. Khandual. J. Indian Chem. Soc. 67, 561 (1990).
15. K.K. Sengupta, A.K. Chatterjee, and S.P. Moulik. Bull. Chem.
Soc. Jpn. 43, 3841 (1970).
–
possible that up to a certain range of MnO4 concentration
(first slow stage), the active species of the oxidant (i.e., the
protonated form of MnO4– (HMnO4) is reactive) while in the
second stage of the reaction, autocatalysis due to one of the
products (Mn(II)) is operative. The product Mn2+ ions
formed in the first slow stage, after achieving an optimum
16. P.S. Radha Krishnamurthi, R.K. Panda, and J.C. Panigrahi. In-
dian J. Chem. 26A, 124 (1987).
17. H. Singh, J.C. Govil, and S.C. Saksena. J. Indian Chem. Soc.
58, 951 (1981).
–
concentration, start reducing MnO4 ions to Mn3+ and or
–
Mn4+ ions (44), (MnO4 + 3Mn2+ + 8H+ = 3Mn3+ + Mn4+
+
18. U. Mishra, K. Sharma, and V.K. Sharma. J. Indian Chem. Soc.
63, 586 (1986).
4H2O) indicating that the intermediate manganese ions are
the active oxidizing species. On addition of Mn2+, the accel-
eration of rate is due to the existence of Mn3+ or Mn4+ as the
principal reactive species in the second fast stage. But the
oxidation of benzilic acid by Mn(III), as well as the continu-
ous increase in the oxidation rate with increasing Mn2+
19. J.R. Jones, W.A. Waters, and J.S. Littler. J. Chem. Soc. 630
(1961).
20. (a) G.H. Jeffery, J. Bassett, J. Mendham, and R.C. Denney.
Vogel’s textbook of quantitative chemical analysis. 5th ed.
ELBS Longman, Essex, UK. 1996. p. 371; (b) p. 334.
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Dalton Trans. 605 (1988).
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24. N.N. Halligudi, S.M. Desai, and S.T. Nandibewoor. Transition
Met. Chem. (Dordrecht, Neth.), 26, 28 (2001).
25. L.J. Bellamy. The IR spectra of complex organic molecules.
2nd ed. Methuen and Co., London. 1958. p. 425.
26. A.I. Vogel. A textbook of practical organic chemistry includ-
ing qualitative organic analysis. 3rd ed. Longman, Essex UK.
1973. p. 332.
–
added to the MnO4 – benzilic acid reaction mixture, indi-
cates that Mn3+, but not Mn4+, is the sole oxidant throughout
the auto acceleration stage.
Conclusions
In an acidic medium, the reaction exhibits an autocatalytic
nature and follows two stages, as one of the products
(Mn(II)) enhances the rate of the reaction. Mn(II) as the re-
duced product of Mn(VII) in the reaction may suggest that
benzilic acid shows a strong reducing character in a HClO4
medium. The overall mechanistic sequences described in the
27. S.T. Nandibewoor, and V.A. Morab. J. Chem. Soc. Dalton
Trans, 483 (1995).
© 2004 NRC Canada

1380
Can. J. Chem. Vol. 82, 2004
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72, 1439 (1953); S.Bhattacharya and P. Banerjee. Bull. Chem.
Soc. Jpn. 69, 3475 (1996).
[Mn(VII)]T = [Mn(VII)]f + [HMnO4]
= [Mn(VII)]f(1 + K1[H+])
29. S.A. Chimatadar, S.C. Hiremath, and J.R Raju. Indian. J.
Therefore,
Chem. 30A, 190. (1991).
30. K.B. Wiberg. Oxidation in organic chemistry. Part A. Aca-
demic Press, New York. 1965. p. 6.
31. K.B. Wiberg. Oxidation in organic chemistry. Part A. Aca-
demic Press, New York. 1965. p. 57.
32. K.S. Gupta and Y.K. Gupta. J. Chem. Ed. 61, 1972 (1964).
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Comprehensive inorganic chemistry. Vol. 3. Pergamon. Press
Ltd., New York. 1975. p. 771.
–
[A1a] [Mn(VII)]f = [MnO4 ]T/1 + K1[H+]
Similarly,
[A1b] [H+]f = [H+]T/1 + K1[Mn(VII)]
Substituting eqs. [A1a]and [A1b] in eq. [A1]
[A1c] Rate = –d[Mn(VII)]/dt
k1K1[H+ ][BA][Mn(VII)]T
(1 + K1[H+ ]) (1 + K1[Mn(VII)])
34. A. Carrington and M.C.R. Symons. Chem. Rev. 63, 443 (1963).
35. J.M. Malcom and R.M. Noyes. J. Am. Chem. Soc. 74, 2769
(1952).
=
36. G.R. Waterbury, A.M. Hayes, and D.S. Martin. J. Am. Chem.
Soc. 74, 15 (1952).
37. R.S. Verma, M.J. Reddy, and V.R. Shastry. J. Chem. Soc.
Perkin Trans. 2, 469 (1976).
38. S.S. Kim, W. Zhang, and T.J. Pinnavaia. Catal. Lett. 43, 149
(1997).
39. B.Z. Zahn and X.Y. Li. Stud. Surf. Sci. Catal. 105A, 615
(1997).
40. P.S. Radha Krishnamurti and M.D.P. Rao. Indian J. Chem.
The term (1 + K1[Mn(VII)]) in the denominator of eq. [A1c]
approximates unity in view of the low concentration of
Mn(VII) used.
Autocatalysis
According to Scheme 2,
15A, 524 (1977).
41. L.S. Levitte and E.R.J Melinovski. J. Am. Chem. Soc. 77,
4517 (1955).
42. S.R. Sahu, V.R. Chourey, S. Pandey, and V.R. Shastry. Oxid.
Commun. 21, 574 (1998).
43. P.D. Pol, R.T. Mahesh, and S.T. Nandibewoor. J. Chem. Res.
Synop. 533 (s), 1145 (m) (2002).
[A1d] Rate = k2[C2][HMnO4]
= k2K2[Mn(II)]f[BA]f2[HMnO4]
[Mn(II)]T = [Mn(II)]f + [C2]
= [Mn(II)]f(1 + K2[BA]2)
44. W.A. Waters. Q. Rev. Chem. Soc. 12, 277 (1958).
Therefore,
[A1e] [Mn(II)]f = [Mn(II)]T/1 + K2[BA]2
Appendix
According to Scheme 1,
[A1] Rate = k1[HMnO4][BA]
= klK1[Mn(VII)]f[BA]f[H+]f
Similarly,
[A1f] [BA]f = [BA]T/1 + K2[Mn(II)][BA]
Substituting eqs. [A1e] and [A1f] in eq. [A1d]
k2K2[Mn(II)][BA]2[HMnO4]
–
The total concentration of MnO4 is given by (where “T”
[A1g] Rate =
{1 + K2[BA]2}{1 + K2[Mn(II)][BA]}2
and “f “stand for total and free)
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