Oxidation of mandelic acid by alkaline potassium permanganate. A kinetic study

Article abstract of DOI:10.1002/(SICI)1099-1395(199807)11:7<448::AID-POC23>3.0.CO;2-A

The kinetics of the oxidation of mandelic acid (MA) by permanganate in aqueous alkaline medium at a constant ionic strength of 1.0 mol dm^-3 were studied spectrophotometrically. The reaction shows first-order kinetics in [permanganate ion] and f

Full text of DOI:10.1002/(SICI)1099-1395(199807)11:7<448::AID-POC23>3.0.CO;2-A

JOURNAL OF PHYSICAL ORGANIC CHEMISTRY, VOL. 11, 448–454 (1998)
Oxidation of mandelic acid by alkaline potassium
permanganate. A kinetic study
Rafeek G. Panari, Ravindra B. Chougale and Sharanappa T. Nandibewoor*
Post Graduate Department of Studies in Chemistry, Karnatak University, Dharwad 580 003, India
Received 21 May 1997; revised 2 December 1997; accepted 2 December 1997
ABSTRACT: The kinetics of the oxidation of mandelic acid (MA) by permanganate in aqueous alkaline medium at a
constant ionic strength of 1.0 mol dmꢀ3 were studied spectrophotometrically. The reaction shows first-order kinetics
in [permanganate ion] and fractional order dependences in [MA] and [alkali]. Addition of products, manganate and
aldehyde have no significant effect on the reaction rate. An increase in ionic strength and a decrease in dielectric
constant of the medium increase the rate. The oxidation process in alkaline medium under the conditions employed in
the present investigation proceeds first by formation of an alkali permanganate complex, which combines with
mandelic acid to form another complex. The latter decomposes slowly followed by a fast reaction between the free
radical of mandelic acid and another molecule of permanganate to give products. The reaction constants involved in
the mechanism were derived. There is good agreement between the observed and calculated rate constants under
different experimental conditions. The reaction was studied at different temperatures and activation parameters were
computed with respect to the slow step of the proposed mechanism. 1998 John Wiley & Sons, Ltd.
KEYWORDS: mandelic acid; alkaline permanganate; oxidation; kinetics
INTRODUCTION
In strongly alkaline medium, the stable reduction
product9 of permanganate ion is manganate ion,
MnO24ꢀ. No mechanistic information is available which
would permit one to distinguish between a direct one-
electron reduction to Mn (VI) (Scheme 1), or prior
formation of hypomanganate in a two-electron step
followed by a fast reaction (Scheme 2).
Oxidation by permanganate ion has extensive applica-
tions in organic syntheses,1–7 especially since the advent
of phase-transfer catalysis,3,4,6 which permits the use of
solvents such as methylene chloride and benzene. Kinetic
studies are important sources of mechanistic information
on the reactions, as demonstrated by results referring to
unsaturated acids in both aqueous1,3–7 and non-aqueous
media.8
The manganese chemistry involved in these multi-step
redox reactions is an important source of information as
the manganese intermediates are relatively easy to
identify when they have sufficiently long lifetimes and
the oxidation states of the intermediates permit useful
conclusions as to the possible reaction mechanisms,
including the nature of intermediates.
k 30
MnꢁVII† ‡ S ꢀ! MnꢁV† ‡ products
k40
MnꢁVII† ‡ MnꢁV† ꢀ! 2MnꢁVI†
where S = substrate and k40 ꢂ k30
Scheme 2
Mandelic acid is used in the form of its salts as a
bacteriostatic agent for genitourinary tract infections. It
has been oxidized by different oxidants in aqueous
alkaline medium10 and its oxidation by permanganate in
acid medium has been reported.11 Different workers have
reported different products12 for the oxidation of
mandelic acid. Although some work on the oxidation of
organic13 and inorganic14 substrates by permanganate in
aqueous alkaline medium has been carried out, there is no
report in the literature on the oxidation of mandelic acid
in such media. This work was carried out on such a
reaction in order to elucidate the redox chemistry of
permanganate in alkaline media.
k 10
Á
MnꢁVII† ‡ S ꢀ! MnꢁVI† ‡ S
k 20
Á
MnꢁVII† ‡ S ꢀ! MnꢁVI† ‡ products
where S = substrate and k20 ꢂ k10
Scheme 1
*Correspondence to: S. T. Nandibewoor, Post Graduate Department of
Studies in Chemistry, Karnatak University, Dharwad 580 003, India.
1998 John Wiley & Sons, Ltd.
CCC 0894–3230/98/070448–07 $17.50

OXIDATION OF MANDELIC ACID BY ALKALINE PERMANGANATE
449
was checked by preparing the reaction mixture and
following the reaction in an atmosphere of nitrogen. No
significant difference between the results obtained under
nitrogen and in presence of air was observed. In view of
the ubiquitous contamination of basic solutions by
carbonate, the effect of carbonate on the reaction was
also studied. Added carbonate had no effect on the
reaction rate. However, fresh solutions were used when
conducting the experiments.
EXPERIMENTAL
Materials. Stock solutions of mandelic acid (Mallinkrodt)
and potassium permanganate (BDH) were prepared by
dissolving the appropriate amounts of samples in doubly
distilled water. The stock solution of permanganate was
standardized against oxalic acid.15 Potassium manganate
solution was prepared as described by Carrington and
Symons16 as follows: a solution of potassium permanga-
nate was heated to boiling above 120°C in 8.0 mol dmꢀ3
potassium hydroxide solution until a green colour was
produced. The solid potassium manganate formed on
cooling was recrystallized from the same solvent. Using
the required amount of recrystallized sample, a stock
solution of potassium manganate was prepared in
aqueous potassium hydroxide. The solution was standar-
dized by measuring the absorbance using a Hitachi
Model 150-20 spectrophotometer with a 1 cm cell at
608 nm (e = 1530 Æ 20 dm3 molꢀ1 cmꢀ1).
Stoichiometry and product analysis. The reaction
mixture containing excess permanganate over mandelic
acid was mixed in the presence of 0.50 mol dmꢀ3 NaOH
adjusted to a constant ionic strength of 1.0 mol dmꢀ3
.
When the reaction time had elapsed, solid KI was added,
followed by acidification with 10% H2SO4. Then
remaining permanganate was titrated against standard
sodium thiosulphate.13 The results indicated that 2 mol of
MnOꢀ4 consumed 1 mol of mandelic acid according to the
equation
All other reagents were of analytical grade and their
solutions were prepared by dissolving the requisite
amounts of the samples in doubly distilled conductivity
water. NaOH and NaClO4 were used to provide the
required alkalinity and to maintain the ionic strength,
respectively.
C6H5CH(OH)COOH ‡ 2MnOꢀ4 ‡ 2OHꢀ
ꢀ! C6H5CHO ‡ CO2 ‡ 2MnO24ꢀ ‡ 2H2O
ꢁ1†
The main oxidation products were identified as an
aldehyde17 (by a spot test) and manganate. The presence
of benzaldehyde as an oxidation product was also
confirmed by preparing its 2,4-dinitrophenylhydrazone
derivative and comparing its melting point with that of an
authentic sample. Such products were also obtained in
previous work.18
It was further observed that the aldehyde does not
undergo further oxidation under the present kinetic
conditions.
Kinetic measurements. All kinetic measurements were
performed under pseudo-first-order conditions with
mandelic acid in at least a 10-fold excess over
permanganate ion at a constant ionic strength of 1.0
mol dmꢀ3. The reaction was initiated by mixing
previously thermostated solutions of MnOꢀ4 and mandelic
acid which also contained the necessary quantities of
NaOH and NaClO4 to maintain the required alkalinity
and ionic strength, respectively. The temperature was
maintained at 26 Æ 0.1°C. The course of the reaction was
followed by monitoring the decrease in the absorbance of
MnOꢀ4 in a 1 cm quartz cell of a Hitachi model 150-20
spectrophotometer at its absorption maximum of 526 nm
as a function of time. Earlier it was verified that there is
negligible interference from other reagents at this
wavelength. The application of Beer’s law to permanga-
nate at 526 nm had earlier been verified, giving
e = 2083 Æ 50 dm3 molꢀ1 cmꢀ1 (literature e = 2200 dm3
molꢀ1 cmꢀ1). The first-order rate constants k(obs) were
evaluated by plots of log[permanganate] versus time. The
first-order plots in almost all cases were linear up to 80%
of the reaction and the k(obs) values were reproducible to
within Æ5%.
RESULTS
Reaction order
The reaction orders were determined from the slopes of
log k(obs) versus log(concentration) plots by varying the
concentration of oxidant, reductant and alkali in turn,
while keeping the others constant.
Effect of oxidant and substrate
The potassium permanganate concentration was varied in
the range 9.0 Â 10ꢀ5–9.0 Â 10ꢀ4 mol dmꢀ3 and the
linearity of plots of log[MnOꢀ4 ] versus time indicated a
reaction order in [MnOꢀ4 ] of unity. This was also
confirmed by varying the [MnOꢀ4 ], which did not show
any change in pseudo-first-order rate constants k(obs)
(Table 1). The substrate, mandelic acid concentration was
varied in the range 9.0 Â 10ꢀ4–9.0 Â 10ꢀ3 mol dmꢀ3 at
26°C keeping all other reactant concentrations and
In the course of the measurements, the colour of the
solution changed from violet to blue and further to green.
The spectrum of the green solution was identical with that
of MnO24ꢀ. It is probable that the blue colour originated
from the violet of permanganate and the green from
manganate, excluding the accumulation of hypomanga-
nate.
The effect of dissolved oxygen on the rate of reaction
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450
R. G. PANARI, R. B. CHOUGALE AND S. T. NANDIBEWOOR
Table 1. Effect of [MA], [MnOꢀ4 ] and [OHꢀ] on the oxidation of mandelic acid by
permanganate in aqueous alkaline medium at 26°C (I = 1.0 mol dmꢀ3
)
k(obs) Â 103 (sꢀ1
)
[MA] Â 103
[MnOꢀ4 Â 104
[OHꢀ]
(mol dmꢀ3
)
(mol dmꢀ3
)
(mol dmꢀ3
)
Experimental
Calculated
0.9
2.0
3.0
5.0
9.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
0.9
2.0
3.0
7.0
9.0
3.0
3.0
3.0
3.0
3.0
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.1
0.3
0.5
0.7
1.0
0.90
1.74
2.25
3.24
5.81
2.20
2.25
2.23
2.26
2.22
0.85
1.70
2.25
2.70
3.15
0.86
1.71
2.34
3.33
4.63
2.34
2.34
2.34
2.34
2.34
0.89
1.60
2.34
2.66
2.85
conditions constant (Table 1). The reaction order in
[mandelic acid] was found to be less than unity (Table 1).
medium. The plot of logk(obs) versus 1/eT was linear (Fig.
1).
Effect of alkali
Effect of initially added products
The effect of alkali on the reaction was studied at
constant concentrations of mandelic acid and potassium
permanganate and keeping a constant ionic strength of
1.0 mol dmꢀ3 at 26°C. The rate constants increased with
increase in [alkali] (Table 1).
The initially added products such as manganate and
aldehyde did not have any significant effect on the rate of
the reaction.
Test for free radicals
Effect of ionic strength
The reaction mixture was kept for 1 h with acrylonitrile
scavenger in an inert atmosphere. On diluting with
methanol, the formation of a precipitate indicates free
radical intervention in the reaction.
The effect of ionic strength was studied by varying the
sodium perchlorate concentration in the reaction med-
ium. The ionic strength of the reaction medium was
varied from 0.5 to 2.0 mol dmꢀ3 at constant concentra-
tions of permanganate, mandelic acid and alkali. It was
found that the rate constant increased with increasing
concentration of NaClO4 and the plot of logk(obs) versus
I1/2 was linear with a positive slope (Fig. 1).
Effect of temperature
The rate constants, k, of the slow step of Scheme 3 were
obtained from the intercepts of the plots of 1/k(obs) versus
1/[MA] at different temperatures and used to calculate
the activation parameters (Table 2). The values of k were
(0.90 Æ 0.04) Â 10ꢀ2, (1.67 Æ 0.08) Â 10ꢀ2 and (2.85 Æ
0.14) Â 10ꢀ2 sꢀ1 at 26, 31 and 36°C, respectively.
Effect of solvent polarity
The relative permittivity (eT) effect was studied by
varying the tert-butyl alcohol–water content in the
reaction mixture with all other conditions being main-
tained constant. Attempts to measure the relative
permittivities failed. However, they were computed from
the values of pure liquids as in earlier work.19 There was
no reaction of the solvent with the oxidant under the
experimental conditions used. The rate constant, k(obs)'
increased with decrease in the dielectric constant of the
DISCUSSION
The permanganate ion, MnOꢀ4 , is a powerful oxidant in
aqueous alkaline medium. As it exhibits a multitude of
oxidation states, the stoichiometric results and pH of the
reaction media play an important role. Under the present
experimental conditions at pH > 12, the reduction
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OXIDATION OF MANDELIC ACID BY ALKALINE PERMANGANATE
451
Figure 1. Plots of log k(obs) versus I1/2 and versus 1/eT
product of Mn(VII), i.e. Mn(VI), is stable and further
reduction of Mn(VI) might be stopped.14 Diode-array
rapid scan spectrophotometric studies have shown that at
pH > 12, the product of Mn(VII) is Mn(VI) and no
further reduction was observed as reported by Simandi et
al.14a However, on prolonged standing Mn(VI) is slowly
reduced to Mn(IV) under our experimental conditions.
The reaction between mandelic acid and permanganate
in alkaline medium has a stoichiometry of 1:2 with a
fractional order dependence on both alkali and mandelic
acid concentrations and a first-order dependence on
permanganate concentration. No effect of the products
was observed. The results suggest that first the alkali
combines with permanganate to give an alkali–
permanganate complex,20 which then reacts with the
substrate, mandelic acid, to give another complex. The
latter complex decomposes in a slow step to give a free
radical derived from decarboxylated mandelic acid,
which further reacts with another molecule of permanga-
nate in a fast step to yield the products (Scheme 3).
The probable structure of the complex (C2) is
Table 2. Thermodynamic activation parameters for the
oxidation of mandelic acid by alkaline permanganate
activation parameters with respect to the slow step of
Scheme 3
Parameter
Value
Attempts to obtain the UV–visible spectral evidence
for the complex formation failed. However, the interac-
tion might be feeble and such complex formation be-
tween a substrate and an oxidant have been observed in
other studies.21 The formation of the complex is proved
kinetically, i.e. by the non-zero intercept of the plot of
Ea
65.5 Æ 3.5 kJ molꢀ1
LogA
DS≠
DH≠
DG≠
9.5 Æ 0.5
ꢀ72 Æ 3 J Kꢀ1 molꢀ1
63 Æ 3 kJ molꢀ1
85.5 Æ 4.0 kJ molꢀ1
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JOURNAL OF PHYSICAL ORGANIC CHEMISTRY, VOL. 11, 448–454 (1998)

452
R. G. PANARI, R. B. CHOUGALE AND S. T. NANDIBEWOOR
Scheme 3
1/k(obs) vs 1/[MA]. Since Scheme 3 is in accordance with
the generally well accepted principle of non-comple-
mentary oxidations taking place in sequences of one-
electron steps, the reaction between the substrate and
oxidant would afford a radical intermediate. A free
radical scavenging experiment revealed such a possibility
(see below). This type of radical intermediate has also
been observed in earlier work22 on the alkaline
permanganate oxidation of various organic substrates.
Scheme 3 leads to the following rate law:
The terms (1 ‡ K1 K2 [MnOꢀ4 ] [OHꢀ]) and (1 ‡ K1
[MnOꢀ4 ] ‡ K1 K2 [MA] [MnOꢀ4 ]) in the denominator of
Eqn (2) approximate to unity in view of the low
concentration of MnOꢀ4 used. Therefore, Eqn (2)
becomes
d‰MnOꢀ4 Š
dt
kK1K2‰MAŠ‰MnOꢀ4 Š‰OHꢀŠ
1 ‡ K1‰OHꢀŠ ‡ K1K2‰MAŠ ‰OHꢀŠ
ꢁ3†
rate ˆ ꢀ
ˆ
d‰MnOꢀ4 Š
dt
kK1K2‰MAŠ‰MnOꢀ4 Š‰OHꢀŠ
rate ˆ ꢀ
ˆ
ꢁ1 ‡ K1‰OHꢀŠ ‡ K1K2‰MAŠ‰OHꢀŠ†
ꢁ2†
1
Â
ꢁ1 ‡ K1K2‰MnOꢀ4 Š‰OHꢀŠ†ꢁ1 ‡ K1‰MnOꢀ4 Š ‡ K1K2‰MAŠ‰MnOꢀ4 Š†
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JOURNAL OF PHYSICAL ORGANIC CHEMISTRY, VOL. 11, 448–454 (1998)

OXIDATION OF MANDELIC ACID BY ALKALINE PERMANGANATE
453
Figure 2. Plots of 1/k(obs) versus 1/[MA] and versus 1/[OHꢀ]. Conditions as in Table 1
or
The effect of increasing ionic strength on the rate
qualitatively explains the reaction between the same
charged ions as shown in Scheme 3. The effect of solvent
on the reaction kinetics has been described in detail in the
older literature.23–28 For the limiting case of a zero angle
approach between two dipoles or an ion–dipole system,
Amis27 has shown that a plot log k(obs) versus 1/eT gives a
straight line with a negative slope for a reaction between
a negative ion and a dipole or two dipoles, and with a
positive slope for a positive ion and dipole interaction. In
the present study, an increase in rate with decrease in the
dielectric constant of the medium was observed, which
cannot be explained by Amis theory,27 as the presence of
a positive ion is unlikely in the alkaline medium
employed. Applying the Born equation, Laidler and
Eyring derived the equation
rate
kK1K2‰MAŠ ‰OHꢀŠ
ˆ k
ˆ
ꢁobs†
‰MnOꢀ4 Š
1 ‡ K1‰OHꢀŠ ‡ K1K2‰MAŠ ‰OHꢀŠ
ꢁ4†
Equation (4) can be rearranged to the following form,
which is used for verification of the rate law:
1
1
1
1
ˆ
‡
‡
ꢁ5†
k
kK1K2 ‰MAŠ ‰OHꢀŠ kK2 ‰MAŠ
k
ꢁobs†
According to Eqn (5), the plots of 1/k(obs) vs 1/[MA]
and 1/k(obs) vs 1/[OHꢀ] should be linear, as verified in
Fig. 2. The slopes and intercepts of such plots lead to
values of k, K1 and K2 at 26°C of (0.9 Æ 0.04) Â 10ꢀ2
s
ꢀ1, 1.6 Æ 0.08 dm3 molꢀ1 and 265 Æ 13 dm3 molꢀ1
,
respectively. Using these values, rate constants under
different experimental conditions were calculated and
compared with experimental data (Table 1). There is
reasonable agreement between them, which supports
Scheme 3.
ꢀ
ꢁ
N Z2 e2
1
1
rÃ
ln k ˆ ln k0 ‡
ꢀ
ꢁ6†
2 D R T
r
where k0 is the rate constant in a medium of infinite
1998 John Wiley & Sons, Ltd.
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454
R. G. PANARI, R. B. CHOUGALE AND S. T. NANDIBEWOOR
4. D. G. Lee. in Oxidation in Organic Chemistry, edited by W. S.
Trahanovsky, Part D. p. 147, Academic Press, New York (1982).
5. L. I. Simandi. in The Chemistry of Functional Groups, edited by S.
Patai and Z. Rappoport, Suppl. C, Chapt. 13. Wiley, Chichester
(1983).
6. D. G. Lee, E. G. Lee and K. C. Brown. Phase Transfer Catalysis,
New Chemistry, Catalysts and Applications, ACS Symposium
Series No. 326, p. 82. American Chemical Society, Washington,
DC (1987).
dielectric constant and r and r* refer to the radius of the
reacting species and activated complex, respectively. It
can be seen from Eqn (6) that the rate should be greater in
a medium of lower dielectric constant when r* > r. There
is a possibility of intramolecular hydrogen bonding that
could stabilize the transition state, increasing the size of
activated complex by attracting solvent molecules due to
solvation effect. The fairly high positive values of DH≠
and DG≠ (Table 2) also indicate that the transition state is
highly solvated, which results in an increase in the size of
transition state. It is likely that r* > r for mandelic acid,
thus explaining the experimental observations. Hence
one can expect intramolecular hydrogen bonding in
mandelic acid since it contains OH and COOH groups on
the same carbon atom. Such hydrogen bonding is
common for molecules having COOH and NH2, COOH
and OH groups, etc., either on the adjacent carbon atom
or on the same carbon atom of the molecule as found29 in
simple amino acids.
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favourable for electron transfer processes. The value of
DH≠ was due to release of energy of solution changes in
the transition state. The negative values of DS≠ within the
range for radical reactions have been ascribed30 to the
nature of electron pairing and electron unpairing
processes, and to the loss of degrees of freedom, formerly
available to the reactions on the formation of a rigid
transition state. It is also interesting that the oxidant
species [MnOꢀ4 ] required a pH > 12, below which the
system becomes disturbed and the reaction will proceed
further to give a reduced product of the oxidant as
Mn(IV), which slowly develops yellow turbidity. Hence
it becomes apparent that in carrying out this reaction the
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noteworthy that under the conditions studied, the reaction
occurs in two successive one-electron reductions
(Scheme 3) rather than two-electron reduction in a single
step (Scheme 2).
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Acknowledgement
The authors thank Dr S. M. Tuwar of Karnatak Science
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