Kinetics and Mechanism of Oxidation of L-Proline by Heptavalent Manganese: A Free Radical Intervention and Decarboxylation

Abstract

The kinetics of oxidation of L-proline by permanganate in alkaline medium was studied spectrophotometrically. The reaction is first order with respect to[MnO4-] and is an apparent less than unit order, each in [L-proline] and [alkali] under the experimental conditions. The reaction rate increases with increase in ionic strength and decrease in solvent polarity of the medium. Addition of reaction products has no effect on the reaction rate. A mechanism involving the formation of a complex between the oxidant and substrate has been proposed. The reaction constants involved in the mechanism were evaluated. There is a good agreement between the observed and calculated rate constants under varying experimental conditions. The activation parameters with respect to the slow step of the proposed reaction scheme were evaluated and discussed.

Full text

PAPER: 04/2353
* Correspondent.
Kinetics and mechanism of oxidation of L-proline by heptavalent
manganese: a free radical intervention and decarboxylation
R.T. Mahesh, M.B. Bellakki and S.T. Nandibewoor*
P.G. Department of Studies in Chemistry, Karnatak University, Dharwad-580 003, India
The kinetics of oxidation of L-proline by permanganate in alkaline medium was studied spectrophotometrically.
The reaction is first order with respect to[MnO4-] and is an apparent less than unit order, each in [L-proline] and
[alkali] under the experimental conditions. The reaction rate increases with increase in ionic strength and decrease
in solvent polarity of the medium. Addition of reaction products has no effect on the reaction rate. A mechanism
involving the formation of a complex between the oxidant and substrate has been proposed. The reaction constants
involved in the mechanism were evaluated. There is a good agreement between the observed and calculated rate
constants under varying experimental conditions. The activation parameters with respect to the slow step of the
proposed reaction scheme were evaluated and discussed.
Keywords: kinetics, L-proline, oxidation, Mn(VII)
Permanganate ion oxidises a great variety of substrates and
finds extensive applications in organic syntheses,1,2 especially
after the advent of phase transfer catalysis3,4 which permits
the use of solvents such as methylene chloride and benzene.
Kinetic studies constitute an important source of mechanistic
information on the reaction, as demonstrated by results
referring to unsaturated acids in both aqueous1,3 and non-
aqueous media.5
During oxidation by permanganate, it is evident that the
Mn(VII) in permanganate is reduced to various oxidation
states in acidic, alkaline and neutral media. Furthermore,
the mechanism by which this multivalent oxidant oxidises a
substrate depends not only on the substrate but also on the
medium6 used for the study. In strongly alkaline medium,
the stable reduction product7 is the manganate ion, MnO42-.
No mechanistic information is available to distinguish
between a direct one-electron reduction to Mn(VI) and a
mechanism in which a hypomanganate is formed in a two-
electron step followed by a rapid reaction.7 Amino acids have
been oxidised by a variety of oxidising agents.8 The oxidation
of amino acids is of interest as the oxidation products differ
for different oxidants.9,10
L-proline is a non-essential amino acid and is an important
constituent of collagen. As per a recent report,11 L-proline is
considered to be the world’s smallest natural enzyme and it is
used in catalysing the aldol condensation of acetone to various
aldehydes with high stereo-specificity. Although some work
on the oxidation of organic12 and inorganic13 substrates by
permanganate in aqueous alkaline medium has been carried
out, there has been no report on the oxidation of L-proline
in such media. Here the title reaction has been carried out in
order to elucidate the redox chemistry of permanganate in
alkaline media and to arrive at a plausible mechanism.
Experimental
Materials and reagents: Stock solutions of L-proline (sd-fine chemicals)
and KMnO4 (BDH) were prepared by dissolving the appropriate
amounts of samples in double distilled water. The preparations of
other stock solutions are as described previously.7
Kinetic measurements: 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 kinetic
accessory (HI-TECH SFA-12). All kinetic measurements and other
experiments are as given in an earlier paper.7
Results
Stoichiometry and product analysis: The reaction mixtures containing
an excess of permanganate over L-proline, NaOH and adjusted ionic
strength of 0.50 mol dm-3, was allowed to react for 2 h at 25 ± 0.1° C.
The remaining permanganate was then analysed spectrophotometri-
cally. The results indicated that two moles of MnO4- were consumed
by one mole of L-proline as given by Eqn (1).
The products were eluted with solvent ether and the organic
product was submitted to spot tests. The main reaction products
were identified as aminobutaraldehyde the by spot test14 for amine
and aldehyde groups and manganate by its UV-Vis spectrum. The
presence of aldehyde was also confirmed by IR spectroscopy,15
which showed bands at 3444 cm-1 for–NH stretching, 1733 cm-1 for
aldehydic-CO stretching, and 2919 cm-1 for aldehydic-CH stretching
respectively. The only organic product obtained by oxidation was
aminobutaraldehyde, the yield obtained was found to be about
70% from its 2,4-DNP derivative. It was further observed that the
aldehyde does not undergo further oxidation under prevailing kinetic
conditions. Test for the corresponding acid was negative.
Reaction order: The reaction orders were determined from
the slopes of log kobs vs log (concentration) plots by varying the
concentration of the reductant and OH- in turn while keeping others
constant. The potassium permanganate concentration was varied in
the range of 5.0 × 10-5 – 5.0 × 10-4 mol dm-3 and the linearity of plots
of log [MnO4-] versus time (r>0.9985, s< 0.027) indicated a reaction
order of unity in [MnO4-]. This was also confirmed by variation of
[MnO4-], which did not result any change in the pseudo first-order
rate constants, kobs (Table 1). The substrate, L-proline concentration
was varied in the range 5.0 × 10-4 – 5.0 × 10-3 mol dm-3 at 25°C while
keeping all other reactant concentrations and conditions constant. The
reaction order in [L-proline] was found to be less than unity (Table
1). The effect of alkali on the reaction has been studied at constant
concentrations of L-proline and potassium permanganate and a
constant ionic strength of 0.50 mol dm-3. The rate constants increased
with increasing [OH-] (Table 1) (r>0.9978, s<0.018).
Effect of ionic strength and solvent polarity: The effect of ionic
strength was studied by varying the NaClO4 concentration in
the reaction medium. The ionic strength was varied from 0.1 to
1.0 mol dm-3 at constant concentrations of permanganate, L-proline
and alkali. It was found that the rate constant increased with
increasing concentration of NaClO4; the plots of logkobs versus I1/2
was linear with a positive slope (Fig. 1) (r>0.9978, s <0.018).
The effect of relative permitivity (εT) on the rate constant was
studied by varying the t-butanol-water content in the reaction mixture
with all other conditions being maintained constant. Attempts to
measure the relative permitivities were not successful. However, they
were computed from the values of the pure liquids.16 No reaction
of the solvent with the oxidant occurred under the experimental
conditions employed. The rate constant, kobs increased with decreasing
dielectric constant of the medium. The plot of logkobs versus.1/ εT was
linear with a positive slope (Fig.1).
Effect of initially-added products: Initially-added reaction
products such as manganate and aminobutaraldehyde did not show
any significant effect on the rate of the reaction.
NH–CH2–CH2–CH2–CH–COOH + 2MnO4-+2 OH-
→ H2N–CH2–CH2–CH2–CHO + 2MnO42- + H2O +CO2
(1)
JOURNAL OF CHEMICAL RESEARCH 2005 JANUARY, 13–17 RESEARCH PAPER 13
* Correspondent. E-mail: stnandibewoor@yahoo.com

PAPER: 04/2353
14 JOURNAL OF CHEMICAL RESEARCH 2005
Effect of temperature: The rate constants, k, of the slow step of the
mechanism were obtained from the plots of 1/kobs versus 1/[proline] at
four different temperatures. The values of k at different temperatures
are given in Table 2. From the plot of log k versus 1/T (r ≥0.999 and
s < 0.0063) with least square analysis, the activation parameters are
calculated and are given in Table 2.
The thermodynamic parameters of the first step of Scheme 1 and the
activation parameters for the rate-determining step of Scheme 1 could
be evaluated as follows: the hydroxyl ion concentration was varied at
several temperatures and values of K1 were determined. The values
of K1 obtained are 3.80, 4.25, 4.80 and 5.40 dm3 mol-1 at 25, 30, 35
and 40°C, respectively. A vant Hoff’s plot was drawn for the variation
of K1 with temperature (log K1 versus 1/T; r>0.9996, s<0.0173),
and values for ∆H, ∆S and ∆G of 18.6±1.0 kJ mol-1, 72.4±2 JK-1
mol-1, and-3.3±0.2 kJ mol-1 were derived. An Arrhenius plot of log
k versus 1/T (r>0.9996, s <0.0173) yielded the activation parameters
for the rate limiting step of Scheme 1 which are given in Table 2.
A comparison of the values with those obtained for the slow step of
the reaction shows that these values mainly refer to the rate limiting
step, supporting the fact that the reaction before the rate determining
step is fairly rapid and involves little activation energy.18
Discussion
The permanganate ion, MnO4- is a powerful oxidant in aqueous
alkaline medium. As it exhibits multitude oxidation states, the
stoichiometric results and pH of reaction media play a significant
role. At pH>12, the reduction product of Mn(VI) might be stopped.7
The diode array rapid scan spectrophotometer (DARSS) studies
have show that at pH>12 the product of the reaction of Mn(VII) is
Mn(VI) and no further reduction was observed as reported by earlier
work.13 However, on long standing for about 16 h, Mn (VI) is slowly
reduced to Mn(IV) under our experimental conditions.
The reaction between L-proline and permanganate in alkaline
medium has stoichiometry of 1:2 with unit order in [permanganate]
and less than unity in [L-proline] and [OH-]. No products effect
was observed. It is known that L-proline exists in the form of
Table 1 Effect of variation of [MnO4-], [L-proline], and [OH-] on oxidation of L-proline by alkaline permanganate at 25°C,
I = 0.20 mol dm-3
[MnO4-] x 104 [L-proline] × 103 [OH-] kobs × 102 kcal × 102
/mol dm-3 /mol dm-3 /mol dm-3 /s-1 /s-1
0.5 2.0 0.1 7.62 7.60
1.0 2.0 0.1 7.68 7.60
2.0 2.0 0.1 7.60 7.60
4.0 2.0 0.1 7.59 7.60
5.0 2.0 0.1 7.64 7.60
2.0 0.5 0.1 3.90 3.7
2.0 1.0 0.1 5.60 5.4
2.0 2.0 0.1 7.60 7.6
2.0 4.0 0.1 8.60 8.4
2.0 5.0 0.1 9.20 8.7
2.0 2.0 0.02 3.45 3.74
2.0 2.0 0.04 5.29 5.31
2.0 2.0 0.08 7.13 7.10
2.0 2.0 0.10 7.60 7.60
2.0 2.0 0.20 9.68 9.57
0.30
0.50
0.70
0.90
1.10
1.30
1.50
0.20 0.40 0.60 0.80 1.00 1.20 1. 40
3+ log k
obs
I
1/2
2+ log k
obs
0.00
0.04
0.08
0.12
0.16
1.20 1.40 1. 60 1.80
1/ε
T
× 10
2
Polymerisation study: The reaction mixture was mixed with acryloni-
trile monomer and kept for 3 h in a dinitrogen atmosphere. On
diluting the reaction mixture with methanol, a white precipitate was
formed, indicating the intervention of free radicals in the reaction.
The blank experiments of either permanganate or L-proline with
acrylonitrile alone did not induce polymerisation under the same
conditions as those induced with reaction mixtures. Initially-added
acrylonitrile decreases the rate, indicating free radical intervention,
which was the case in earlier work.17
Fig. 1 Plots of (a) log kobs versus I1/2 and (b) log kobs versus
1/εT.
(a) [MnO4-] = 2.0 × 10-4 , [L-proline] = 2.0 × 10-3, [OH-] = 0.10;
(b) [MnO4-] = 2.0 × 10-4, [L-proline] = 2.0 × 10-3,[OH-] = 0.10,
[I] = 0.20 mol dm-3.
Table 2 Activation parameters for the oxidation of L-proline
by permanganate in aqueous alkaline medium with respect to
slow step of Scheme 1
(a) Effect of temperature
Temperature /°C k x 102 /s-1
25 1.03
30 1.72
35 3.20
40 4.40
(b) Activation parameters
Ea ∆H# ∆S# log A ∆G#
/kJ mol-1 /kJ mol-1 /JK-1 mol-1 /kJ mol-1
76 ± 2 74 ± 2 –34 ± 1 11 ± 1 84 ± 4

PAPER: 04/2353
JOURNAL OF CHEMICAL RESEARCH 2005 15
a Zwitter ion19 in aqueous medium. In highly acidic medium,
it exists in the protonated form, whereas in highly basic medium,
it is in the fully deprotonated form.
NH–CH2–CH2–CH2–CH–COO-
The observed fractional order in [OH-] indicates that first alkali
combines with permanganate to form an alkali-permanganate species
[MnO4.OH]2- in a pre-equilibrium step,20 which is also supported by the
Michaelis–Menton plot (Fig 3) which is linear with a positive intercept.
L-proline in the deprotonated form reacts with alkali-permanganate
species [MnO4.OH]2- to form a complex (C). C decomposes in a
slow step to give a free radical derived from decarboxylated L-proline.
Mn
MnO
4-
+ OH
-
K
1
O
O
OOH
O
2-
2
H
•
•
•
N
2
CH
CH
CH
2
CH COO
-
+
Mn
K
2
O
O
OOH
O
2-
Complex (C)
Complex (C) k
slow
+ HCO
3
-
+ MnO
4
2-
2
2
CH
CH CH
CH
2
N
H
Mn
O
CH2
CH2
CH2
CH
N
H
+
O
O
O
OH
H
2N–H2C–H2C–CH2CHO+MnO4
2-
2-
Scheme 1
O
2
Mn
2
2
N
CH
OO
O
O
CH
CH
CH
H
CO
3-
OH
The probable structure of complex (C) is
0.0
0.5
1.0
1.5
2.0
2.5
3.0
3.5
0204060
0.0
0.5
1.0
1.5
2.0
2.5
0.00.51.01.52.02.5
1/[OH
-
] dm
3
mol
-1
1/k
obs
× 10
2
s
1/[L-proline] × 10
-3
dm
3
mol
-1
1/k
obs
× 10
2
s
Fig. 2 Verification of rate law(8): Plots of 1/kobs versus 1/
[L-proline] and1/kobs versus 1/[OH-] (conditions as in Table 1).
This radical in turn reacts with another molecule of permanganate
species in a fast step to yield the products (Scheme 1).
Spectral evidence for complex formation between oxidant and
substrate was obtained from the UV-Vis spectra of the alkali
permanganate species and L-proline. A bathochromic shift, λmax, of ca
6nm from 304 to 310nm is observed, together with hyperchromicity
at λmax 310 nm. Analogous effects upon complex formation between a
substrate and an oxidant have been observed in other investigations.20
Furthermore, the formation of the complex is proved kinetically
by the non zero intercept of the plot of 1/kobs vs 1/[proline] (Fig.2)
(r>0.9935, s<0.043). The observed modest enthalpy of activation
and a relatively low value of the entropy of activation indicate
that oxidation presumably occurs by an inner-sphere mechanism.
This conclusion is supported by earlier work.21 Since Scheme 1 is
in accordance with the generally well-accepted principle of non-
complementary oxidations taking place in sequences of one-electron
steps, the reaction would involve a radical intermediate. A free radical
scavenging experiment revealed such a possibility (vide infra).
2+ log k
1
at 298 K
3
2
5
4
1
-0.5
-0.3
-0.1
0.1
0.3
0.5
0.7
0.9
1.1
1.3
1.5
-0.5 0.5 1.5 2.5
3+ log k
2
at 303 K
Fig. 3 Plot of log k2 at 303 K versus log k1 at 298 K for isokinetic
temperature. (conditions as in Table 3). 1, L-Proline; 2, L (+)
Lysine; 3, L–hydroxy proline; 4, Rac-serine; and 5, L-cystine.

PAPER: 04/2353
16 JOURNAL OF CHEMICAL RESEARCH 2005
of earlier studies on other amino acids. In case of oxidation of
L-Asparagine–permanganate reaction,23 there was no effect of ionic
strength and relative permitivity on the rate of reaction. In case of
L-phenylalanine24 also the rate was independent of ionic strength.
But in our present study the effect of increasing ionic strength on
the rate explains qualitatively the reaction between two negatively
charged ions, as seen in Scheme 1. However, increasing the content
of t-butyl alcohol in the reaction medium leads to the increase in the
reaction rate, contrary to the expected slower reaction between like
ions in the media of lower relative permitivity. Perhaps the effect is
countered substantially by the formation of active reactive species to
a greater extent in low relative permitivity media leading to the net
increase in rate.25 The entropy of activation (∆S#) in earlier work23
on L-Asparagine was found to be positive indicating the complex is
less ordered than reactant molecules, whereas in case24 of L-phenyl-
alanine, negative value of ∆S# and high value of frequency factor
indicates that the electrostatic effects are un important. However, in
the present study, the negative value of ∆S# indicates that the complex
is more ordered than the reactants. The activation parameters for
the oxidation of some amino acids by alkaline permanganate are
summarised in Table 3. The entropy of the activation for the title
reaction falls within the observed range. Variation in the rate within
a reaction series may be caused by change in the enthalpy or entropy
of activation. Changes in the rate are caused by changes in both ∆H#
and ∆S#, but these quantities vary extensively in a parallel fashion.
A plot of ∆H# versus ∆S# is linear according to equation,
∆H# = β∆S# + constant
β is called the isokinetic temperature; it has been asserted that
apparently linear correlation of ∆H# with ∆S# are sometimes
misleading and the evaluation of β by means of the above equation
lacks statistical validity.26 Exner27 advocates an alternative method
for the treatment of experimental data. If the rates of several reactions
in a series have been measured at two temperatures and log k2 (at T2)
is linearly related to log k1 (at T1) i.e. log k2 = a + b logk1, he proposes
that β can be evaluated from the equation,
β = T1T2 (b-1)/T2b – T1
We have calculated the isokinetic temperature as 235 K by plotting
log k2 at 303 K vs log k1 at 298 K (Fig.3) (r ≥ 0.998 and s < 0.0067).
The value of β (235 K) is lower than experimental temperature
(298 K). This indicates that the rate is governed by the entropy
of activation.28 The linearity and the slope of the plot obtained
may confirm that the kinetics of these reactions follow a similar
mechanism, as previously suggested.
Conclusion
It is interesting to note that the oxidant species [MnO4-] 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 in the reaction solution.
Hence, it is apparent that in carrying out this reaction the role of pH
in a reaction medium is crucial. It is also noteworthy that under the
conditions studied the reaction occurs in two successive one-electron
reductions rather than a single two-electron step. The rate constant
of the slow step and other equilibrium constants involved in the
mechanism have been evaluated and activation parameters with
respect to the slow step of the reaction were computed.
Received 18 February 2004; accepted 13 September 2004
Paper 04/2353
This type of radical intermediate has also been observed in earlier
work22 on alkaline-permanganate oxidations of amino acids.
Scheme 1 leads to the rate law as follows,
Rate = k [C] = kK1 K2 [L-proline]f [MnO4-]f [OH]f. (2)
Now,
[MnO4-]t = [MnO4-]f +[MnO4 .OH]2- +[C]
=[MnO4-]f (1 + K1 [OH- ] +K1K2[L-proline][OH-]) (3)
Where t and f stand for total and free,
[MnO4-]t
[MnO4-]f = 1 + K1[OH-] + K1K2 [L-proline][OH-] (4)
Similarly,
[OH-]t
[OH-]f = 1 + K1[MnO4-] + K1K2 [L-proline][MnO4-] (5)
and
[L-proline]t
[L-proline]f = 1 + K1K2 [MnO4-] [OH-] (6)
Substituting the values of Eqns (4), (5) and (6) in Eqn (2) and
omitting subscripts t and f, we get Eqn (7)
d[MnO4
-
] kK1K2 [L-proline] [OH-][MnO4-]
rate = – dt = (1+K1[OH-] + K1K2 [OH-][L-proline])
1
×
(1+K1K2[MnO4-][OH-])(1+K1[MnO4-]+K1K2[
L
-proline][MnO4-])
(7)
The terms (1+K1K2[MnO4-][OH-]) and (1+K1[MnO-4] + K1K2
[L-proline][MnO-4]) in the denominator of Eqn (7) approximate to
unity in view of the low concentration of MnO-4 used (K1=3.80 and
K2 =4045). Therefore Eqn (7) becomes
Rate kK1K2 [L-proline] [OH-]
k
obs = [MnO4-] = 1+K1[OH-]+K1K2 [OH-][L-proline] (8)
Eqn (8) can be rearranged to the following form (9), which is suitable
for the verification of the rate law:
1
1 1 1
Kobs
=
kK1K2 [L-proline][OH-] + kK2[L-proline] + k (9)
According to Eqn (9), the plots of 1/kobs versus 1/[L-proline]
(r>0.9988, s<0.045) and 1/kobsvs 1/[OH-](r>0.9928, s<0.052) should
be linear which is verified in Fig. 2. The slopes and intercepts
of the plots lead to the values of k, K1 and K2 at 25 °C of 1.03±
0.04 × 10-2 s-1, 3.80±0.16 dm3 mol-1, and 4.05±0.10 x 103 dm3mol-1,
respectively. Using these values, the rate constants under different
experimental conditions were calculated by Eqn (8) and compared with
experimental data (Table 1). Experimental and calculated values agree
reasonably well, supporting the assumptions of Scheme 1. The value of
K1 is in good agreement with that derived in earlier work.7
The effect of ionic strength and solvent polarity on the rate
of reaction in the present case is entirely different from that
Table 3 Activation parameters for some amino acids (for isokinetic temperature)
Amino acids k x102/dm3 k x102/dm3 ∆S# ∆H# ∆G# Ref.
mol-1s-1 at 298 K mol-1s-1 at 303 K /JK-1mol-1 /kJ mol-1 /kJ mol-1
L-Proline 1.03 1.72 –34 ± 1 74 ± 2 84 ± 4 Present work
L(+) Lysine 0.45 0.68 –28 ± 1 51 ± 2 60 ± 3 29
L-Hydroxy proline 5.0 8.0 –78 ± 4 57 ± 3 81 ± 4 30
Rac-serine 1.0 1.4 –28 ± 2 49 ± 3 58 ± 3 31
L-Cystine 6.1 6.5 –105 ± 10 48.5 ± 2.5 80 ± 4 32

PAPER: 04/2353
JOURNAL OF CHEMICAL RESEARCH 2005 17
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