Kinetics of oxidation of atropine by alkaline KMnO4 in aqueous solutions

Article abstract of DOI:10.14233/ajchem.2015.19039

Kinetics of oxidation of atropine by alkaline KMnO₄ has been studied in aqueous solutions at various temperatures. Kinetic I measurements were performed at constant pH and constant ionic strength, under pseudo-first order condition in which the concentration i of atropine was around an order of magnitude greater than that of KMnO₄. The observed rate constant was obtained by following the change in absorbance of the reaction mixture with time at a predetermined wavelength. The overall rate constant and the orders of the I reaction with respect to concentrations of both atropine and KMnO₄ were determined. Two observed rates and hence, two rates constant | were found. This indicates that the oxidation process occurs in two steps with the formation of intermediates. A suggested mechanism for I the reaction is given which agrees with kinetics results. The activation energy of the process was convoluted from temperature effect on. the rate of the reaction.

Full text of DOI:10.14233/ajchem.2015.19039

Asian Journal of Chemistry; Vol. 27, No. 10 (2015), 3877-3882
A
SIAN
J
OURNAL OF HEMISTRY
C
http://dx.doi.org/10.14233/ajchem.2015.19039
Kinetics of Oxidation of Atropine by Alkaline KMnO4 in Aqueous Solutions
*
ABDULLAH I. SALEH , SANAD K. AL-GHREIZAT and HAMZEH M. ABDEL-HALIM
The Hashemite University, Department of Chemistry, Faculty of Science, P.O. Box 330127, Zarqa 13133, Jordan
*Corresponding author: Fax: +962 5 3903349; Tel: +962 5 3903333; E-mail: a-saleh@hu.edu.jo
Received: 18 January 2015;
Accepted: 22 February 2015;
Published online: 22 June 2015;
AJC-17365
Kinetics of oxidation of atropine by alkaline KMnO4 has been studied in aqueous solutions at various temperatures. Kinetic
measurements were performed at constant pH and constant ionic strength, under pseudo-first order condition in which the concentration
of atropine was around an order of magnitude greater than that of KMnO4. The observed rate constant was obtained by following the
change in absorbance of the reaction mixture with time at a predetermined wavelength. The overall rate constant and the orders of the
reaction with respect to concentrations of both atropine and KMnO4 were determined. Two observed rates and hence, two rates constant
were found. This indicates that the oxidation process occurs in two steps with the formation of intermediates. A suggested mechanism for
the reaction is given which agrees with kinetics results. The activation energy of the process was convoluted from temperature effect on
the rate of the reaction.
Keywords: Kinetics, Atropine, Potassium permanganate, Anticholinergic.
INTRODUCTION
H3C
Atropine sulfate monohydrate (ASM, Fig. 1) is pharmaceu-
N
tical formulation that is one of theWorld Health Organization’s
OH
list of essential medicines1. Atropine itself is a naturally
occurring tropane alkaloid with chemical structure similar to
.H2SO4.H2O
that of cocaine. Plants of the family Solanaceae that includes
Atropa belladonna, Datura stramonium and Mandragora
officinarum are the major natural sources of the alkaloid.
Atropine serves as a drug with a wide variety of effects2. It is
a competitive antagonist of the muscarinic acetylcholine
receptors, hence its major effects are on the parasympathetic
nervous system.As an anticholinergic drug (parasympatholytic)
it is considered a nonselective muscarinic acetylcholinergic
antagonist. Consequently, its main medical uses are based on
these potent biological properties. Since it blocks the action
of the vagus nerve, atropine is used in resuscitation procedures3.
Additionally, atropine mixed with pralidoxime chloride
(2-PAM chloride) is used as an antidote for poisoning by
organophosphate insecticides and nerve gases, such as Tabun
(GA), Sarin (GB), Soman (GD) and VX4. Atropine has also
several medical applications that include ophthalmic uses as
therapeutic mydriatic and in slowing the progression of myopia
O
Ph
O
2
Fig. 1. Chemical structure of atropine sulfate monohydrate
alkaline copper(III) periodate complex has been recently
reported8,9. Organic oxidation protocols have commonly
employed transition metal complexes in their higher oxidation
states that can be stabilized by appropriate polydentate
ligands10,11. Atropine oxidation by typical strong oxidizing
reagents including potassium dichromate and potassium
permanganate in appropriate pH media have never been
reported to the best of our knowledge. Understanding the
behaviour of atropine when treated with powerful oxidizing
agents in aqueous media might shed some light into its
interaction with oxidative enzymes such as FAD and NADP12.
Kinetics of oxidation of organic substrates by alkaline
KMnO4 has been reported in literature and demonstrated to
be a powerful mechanism elucidating tool. Examples include
in children5, 6
.
As much as it is crucially important to understand this kind
of physiological activity7, atropine interactions with common
chemical reagents are still lacking. Oxidation of atropine by
oxidation of DL-aspartic acid and sugars13,14
.

3878 Saleh et al.
Asian J. Chem.
It is assumed that, under the reaction conditions and
did not proceed to any measurable extent. However, the
addition of MnO2 greatly enhanced the hydrolysis.
stoichio-metries used herein, oxidation of the aliphatic alcohol
of atropine to a carboxyl group is the predominant reaction15
(Scheme-I). Permanganate ion attains its optimum stability
in neutral to slightly alkaline media16. However, in strongly
alkaline media, it disproportionates to manganese(V) (hypo-
manganate) or manganese(VI) (manganate). Hence, for oxida-
tion reactions under these conditions it will be difficult to deter-
mine whether a one or two electron process is operating12. To
sort out the exact reaction of basic KMnO4 with atropine a set
of preliminary experiments were done.
In the present work, as an extension of our previous
studies17-20 on both kinetics and biological application to widen
our knowledge about the oxidation of biologically important
substances, we report the kinetics of oxidation of atropine by
alkaline KMnO4. Kinetic results at various temperatures are
presented and a proposed mechanism for the oxidation process
is given and discussed.
Alkaline KMnO4 disproportionation: A blank run of
mixing all ingredients [KMNO4 (0.001 M), KOH 0.10 M)]
except atropine sulfate monohydrate has shown that the change
in absorption by KMnO4 is negligible throughout the
measuring period.
Alkaline KMnO4 oxidation:Atropine sulfate monohydrate
solution (0.01 M) (about 10 folds in excess) was treated with
the alkaline KMnO4 (0.001 M) solution for 25 min. Products
were isolated, purified and thoroughly characterized. The
products were excess atropine, tropine, tropic acid and phenyl-
malonic acid (Scheme-I). Addition of mercuric chloride or
acrylonitrile monomer to the reaction system, directly after
the appearance of the green colour of MnO42–, induced a white
precipitate formation and polymerization, respectively. Conse-
quently, a radical formation throughout the course of the
reaction mechanism cannot be ruled out21,22. The reaction was
quenched and the green colour persists.
EXPERIMENTAL
These results implicated that the major change in the
oxidant concentration (absorbance at its maximum wavelength)
is attributed to the oxidation of atropine primary alcohol.
Hence, kinetic measurements were designed to follow up this
reaction. It is aimed to investigate the mechanism by which
the reaction is taking place.
Alkaline hydrolysis:Atropine sulfate monohydrate (0.10
M) was treated with KOH (1.0 M) solution for 25 min (optimum
run time frame for reaction monitoring, arbitrary set). Hydrolysis
products were then isolated, purified and characterized by
NMR and IR spectroscopy. Results indicated that hydrolysis
COOH
H C
3
O
N
OH
O
H C
3
O
KMnO4, OH-
1 equivalant
N
OH
+
O
H C
3
O
N
Atropine
~ 10 equivalents
O
Excess Atropine
OH
H C
3
OH
OH
H C
3
N
O
N
Partial Alkaline Hydrolysis
O−
+
O
O
Excess Atropine
Tropic Acid (salt)
Tropine
O-
O
H C
3
OH
COOH
N
H C
O-
3
Alkaline Hydrolysis
O
N
+
O
O
Phenylmalonic acid
(salt)
Tropine
Scheme-I: Main reactions of atropine in presence of KMnO4 in basic medium

Vol. 27, No. 10 (2015)
Kinetics of Oxidation of Atropine by Alkaline KMnO4 in Aqueous Solutions 3879
OH
COOH
Kinetics measurements: Freshly prepared aqueous solutions
of desired concentrations of atropine sulfate monohydrate and
KMnO4 were used in the kinetic study. In a strongly basic
solution (0.1 M), atropine sulfate monohydrate is converted
to atropine. Measurements were carried out using a Diode
Array Spectrophotometer model 8453E from HP Agilent
Technologies. Reactions were monitored by following the
change in absorbance of (MnO42–)reaction mixture with time
at a predetermined wavelength. The wavelength was deter-
mined by recording the absorption spectra for KMnO4 alone
and for its mixture with atropine sulfate monohydrate after
completion of the reaction. The wavelength of maximum absor-
bance difference (λmax), preferably in visible region of the
spectra, between the absorption of KMnO4 and the reaction
mixture was selected. A λ scan for reactants and products at
H3C
H3C
O
O
N
KMnO4, OH-
1 equivalant
N
O
O
Atropine
~ 10 equivalents
An excess of atropine remains. The excess atropine is
expected to undergo partial hydrolysis due to the presence of
MnO2 as byproduct which acts as an auto-catalyst to give
tropine and tropic acid and the oxidized product in stage I is
also expected to undergo an alkaline hydrolysis to yield tropine
and phenylmalonic acid (Scheme-I). The rate of oxidation of
atropine is given by:
Rate = k[ATN]a[KMnO4]b
(1)
various time intervals is shown in Fig. 2, from which λmax
=
where k is the reaction rate constant and a & b are orders of
reaction with respect to concentrations of atropine and KMnO4,
respectively. The concentration of OH– was kept constant at
0.10 M throughout the all kinetics measurements. Under
pseudo-first order conditions in which [ATN] >> [KMnO4],
the concentration of atropine is essentially constant throughout
the reaction. The reaction rate is thus given by:
440 nm was selected. At this wavelength absorbance is mainly
due to MnO42– and absorbance due to all other Mn species is
minimal. Also, it should be noted that none of the organic
compounds, reactants or products, absorb at this wavelength.
1.2
d[KMnO4 ]b
1.0
Rate ≈ −
= kobs[KMnO4 ]b
(2)
(3)
dt
1 min
where kobs is the observed rate for reaction is given by:
kobs = k[ATN] a
2 min
3 min
4 min
5 min
6 min
7 min
8 min
9 min
11 min
MnO4–
0.8
0.6
0.4
0.2
0
where k is the rate constant for eqn. 1.
For a first-order dependence of reaction rate on [KMnO4],
experimental absorbance-time data pairs were fitted to the
exponential function:
(A∞ – At) = (A∞ –A0) exp(–kobs t)






(A∞ − At )
(A∞ − A0 )
ln
= −kobst
or
(4)


where At is the absorbance of the reaction mixture which is
mainly due to MnO42– at a given time (t) through the reaction,
A0 is its initial absorbance (t = 0) and A∞ is the absorbance of
the mixture at the end of the reaction, that is when the absor-
bance no longer changes with time (t = ∞).
200
300
400
500
600
700
800
Wavelength (nm)
Fig. 2. Plot of absorbance vs. λ for the product and for the reactant of the
oxidation of atropine by MnO4–. Kinetic experiment were studied
at λmax = 440 nm
Experimental results showed that a plot of ln [(A∞ – At)/
(A∞ – A0)] vs. time gives a straight line. Its slope represents
the rate of appearance of MnO42–, which equals the rate of
disappearance of MnO4– (as shown in the suggested mechanism
below). The value of kobs (s–1) was obtained from the slope,
according to eqn. 4. Using eqn. 3, a plot of ln kobs vs. ln [ATN]a
is a straight line that gives the reaction rate constant, k, in
units of dm3 mol–1 s–1 (intercept) and the order of the reaction
(a) with respect to [ATN] (slope). At specific concentrations
of substrate and oxidant several trials were carried out for each
measurement and the average value of the observed rate is
reported. Two observed rates were found (kobs1 and kobs2) and
hence, two different reaction rate constants are obtained.
Experimental errors in value of rate constant are estimated to
be 10 %.
All reactions were studied under pseudo-first order
conditions, at which the concentrations of atropine sulfate
monohydrate [10–2 to 10–1 mol dm–3] were at least an order of
magnitude larger than those of KMnO4 [10–4 to 10–2 mol dm–3].
The pH of the reaction mixture was maintained constant at
13.0 0.1 and its ionic strength was fixed using NaClO4. Rates
were measured at various temperatures and for each measure-
ment it was maintained constant within 0.1 °C.
RESULTS AND DISCUSSION
The main reactions of oxidation of atropine by KMnO4
in basic medium are shown in Scheme-I. Since kinetics
measurements were run under pseudo first order condition, in
which the concentration of atropine is an order of magnitude
larger than that for KMnO4, then, the first stage of the process
which involves the oxidation of atropine by alkaline KMnO4
is given by the following equation:
The kinetics results for the oxidation of atropine by
KMnO4 are shown in Table-1 and sample plots are shown in
Figs. 3-6.

3880 Saleh et al.
Asian J. Chem.
TABLE-1
CONCENTRATION EFFECT ON RATE OF OXIDATION OF ATROPINE BY KMnO4 IN AQUEOUS BASIC SOLUTIONS, [KOH] = 0.10 M
[ATN] = 1.0 × 10–2 [MnO4–] = 1.0 × 10–3
M
M
Run
1
2
3
4
[KMnO4] M
2.00 × 10-3
1.75 × 10-3
1.50 × 10-3
1.00 × 10-3
kobs1 × 103
kobs2 × 103
Run
5
6
7
8
[ATN] M
1.75 × 10-2
1.50 × 10-2
1.00 × 10-2
0.50 × 10-2
kobs1 × 103
kobs2 × 103
6.51
6.30
5.40
2.26
2.13
2.18
2.25
5.26
5.64
5.27
5.40
2.80
2.66
2.25
1.91
3.79
Time (s)
900
Based on the obtained kinetics results, we propose the
following mechanism for the oxidation process of atropine.
Step 1
600
700
800
1000
1100
1200
-1.00
-1.20
-1.40
-1.60
-1.80
-2.00
-2.20
-2.40
-2.60
-2.80
-3.00
y = -2.25E-03x - 3.63E-02
R2 = 9.86E-01
K1
−
(5)
ATN
+
[MnO4.OH]2
C
Step 2
k2
2
-
−
C
MnO4
+
I
+ OH
(Slow) (6)
H2O, OH-
Step 3
Step 4
2
−
k3
-
Fig. 3. Sample plot for rate of step 2 of ln [(A∞ – At)/(A∞ – A0)] vs. time;
(7)
MnO4
+
I
II + MnO4
+
H2O
[ATN] = 1.0 × 10–2 M, [KMnO4] = 1.0 × 10–3 M, [KOH] = 0.10 M
K4
Hydrate of II
k5
H2O, OH-
+
[MnO4.OH]2
Time (s)
−
(8)
D
1240
-2.90
1290
1340
1390
1440
Step 5
2-
MnO4
+
product + OH - (Slow)
D
-3.10
-3.30
-3.50
-3.70
-3.90
-4.10
(9)
y = -5.40E-03x + 3.75E+00
R2 = 9.81E-01
Step 6
2
−
(10)
MnO
disproportionates to MnO
2
where C and D are4complexes, I and II are intermediates.
In a strongly basic solution, MnO4– will be in the form of
[MnO4.OH]2–. Experimentally, the reaction rate was followed
by monitoring the change in absorbance of the reaction mixture
with time at λ = 440 nm. At this wavelength, the absorbance
(At) is mainly due to the formation of MnO42– (Fig. 2) and it is
Fig. 4. Sample plot for rate of step 5 of ln [(A∞ - At)/(A∞ –A0)] vs. time;
[ATN] = 1.0 × 10–2 M, [KMnO4] = 1.0 × 10–3 M, [KOH] = 0.10 M
ln [ATN]
–
increasing with time. Due to regeneration of MnO4 from
-5.50 -5.30 -5.10 -4.90 -4.70 -4.50 -4.30 -4.10 -3.90
-5.70
MnO42– (eqn. 7), then the rate of formation of MnO42– is appro-
ximately equal to the rate of disappearance of MnO4–.
Oxidation of primary alcohols under strongly alkaline
KMnO4 is known to occur in two stages23. For the first stage,
an aldehyde intermediate is well-established. Further oxidation
of the aldehyde to the final product (carboxylic acid) proceeds
at a faster rate compared to the first stage.
-5.80
y = 0.3073x - 4.6471
R2 = 0.982
-5.90
-6.00
-6.10
-6.20
-6.30






A∞ − At
A∞ − A0
Fig. 7 depicts the absorbance change ln
over
Fig. 5. Sample plot for rate of step 2 of ln [ATN] vs. ln kobs1; [KMnO4] =


1.0 × 10–3 M, [KOH] = 0.10 M
the whole reaction time. Clearly the reaction is governed by
two distinctive rates. The reason for two stages is based on the
fact that intermediates are reaching their optimum concen-
tration or due to the occurrence of consecutive reactions24.At the
first stage (steps 1-3) of the reaction, the absorbance increased
rapidly with time that indicating that oxidation of atropine to
intermediate I which is relatively slow (compared with stage
2 (steps 4 and 5). The rate limiting step of this stage is
represented by step 2 (eqn. 6). However; at certain time when
an optimum concentration of intermediate II (then the hydrate
form of the aldehyde)25 is reached, the reaction will be
controlled by step 5 (eqn. 9) that is faster than the first limiting
step 3. Finally oxidation of intermediate II proceeds cleanly
to produce the final products.
ln [ATN]
-5.30 -5.10 -4.90 -4.70 -4.50 -4.30 -4.10 -3.90
-4.80
-4.90
y = 0.44x - 3.2282
R2 = 0.9885
-5.00
-5.10
-5.20
-5.30
-5.40
-5.50
-5.60
Fig. 6. Sample plot for rate of step 4 of ln [ATN] vs. ln kobs2; [KMnO4] =
1.0 × 10–3 M, [KOH] = 0.10 M

Vol. 27, No. 10 (2015)
Kinetics of Oxidation of Atropine by Alkaline KMnO4 in Aqueous Solutions 3881
Time (s)
21 141 261 381 501 621 741 861 981 1101 1251 1401
Intermediate I
OH
Complex C
O
0
-1.00
-2.00
-3.00
-4.00
-5.00
-6.00
-7.00
OH
-
H
O
O
C
RO
Mn
-
O
RO
O
O
O
Intermediate II
Complex D
-
O
O
H
O
H
O
O
H
Fig. 7. Plot of ln (absorbance) vs. time for the reaction mixture at λmax = 440
Mn
O
OH
-
nm; [KMnO4] = 1.00 × 10-3 M, [ATN] = 1.00 × 10-2 M, [OH-] = 0.10 M
O
RO
RO
The change in absorbance with time represents the rate
of the reaction according to eqn. 2. From a kinetic point of
view, steps 2 and step 5 will be the rate determining steps and
steps 1, 3, 4, 6 are fast. Experimentally, two rates were observed.
The first reaction rate (rate 1) is due to step 2, consequently,
eqn. 1 above can be rewritten as:
O
O
N
R =
(
)
Rate 1 = k2[C] and from eqn. (5), [C] = K1[ATN][MnO4–]
–
and, [[MnO4.OH]2–] = [MnO4 ], then
Fig. 8. Possible structures of intermediates and complexes involved in the
proposed mechanism
Rate 1 = k2K1[ATN][MnO4–]
Here, k2K1 is a constant quantity which represents the first
rate of the reaction and will be represented as krxn1. Then,
The above proposed mechanism is supported by the obser-
vation of initial colour change of permanganate from violet to
green (manganate). The radical mechanism was invoked due
to the precipitation observed when mercuric chloride was added
to the system28. When HgCl2 was added after the appearance
of the green colour, the colour persists. The reaction was
quenched at this stage.
Rate 1= krxn1[ATN]a[MnO4–]b
(11)
The second reaction rate (rate 2) is due to step 5, given by:
Rate 2 = k5[D]
From eqn. 8, [D] = K4[Hydrate of II][MnO42–] and,
[Hydrate of II] = [II] = [ATNreacted], then
Rate 2 = k5K4[Hydrate of II][MnO42–]
Since the observed overall rate was obtained to be first
order with respect to permanganate, eqns. 11 and 12 above
will have the following expression
2–
The concentration of MnO4 produced equals to the
concentration of reacted MnO4–. The constant quantity k5K4
represent the second rate of the reaction and will be called
krxn2. Then:
Rate = kobs[MnO4–]
(13)
And since [ATN] >> [KMnO4], the observed rate of the
reaction, kobs, is given by:
Rate 2 = krxn2[ATNreacted]a [MnO4–]b
(12)
For step 2: kobs1 = krxn1[ATN]a
A strong reason why a λmax at 440 nm, which is specific
for MnO42–, was selected is due to the formation of MnO42– in
both of the slow steps 2 and 5. The disappearance of the distinct
permanganate absorbance band located in the range of wave-
lengths 490-590 nm and the formation of bands with maxima
at 347, 439 and 606 nm (Fig. 2) are characteristic of manganate
MnO42–, supports this observation. Disproportionation of
manganate to form permanganate and manganese dioxide is
then follows25. Possible structures of the intermediates and
complexes involved are shown in Fig. 8. It is worth mentioning
that the proposed structures of these intermediates are not
conclusive but highly possible26. The aldehyde-intermediate(II)
is logical since it is a common intermediate for similar
oxidations (e.g., K2Cr2O7 oxidation)27. Wiberg and Stewart25
demonstrated the free radical mechanism of oxidation of
aldehydes by alkaline KMnO4 solution. Formation of
permanganate ester of the hydrate of the aldehyde was also
implicated. The regenerated permanganate present in the form
(MnO4.OH)2– will react preferentially with intermediate II
rather than the excess atropine, since it is known that aldehydes
are more reactive compared to primary alcohols toward the
permanganate oxidizing reagent. This step involves a second
single electron transfer.
For step 5: kobs2 = krxn2[ATN]a
(14)
The rates constant krxn1 and krxn2 were obtained from plots
of ln kobs1 vs. ln [ATN] and ln kobs2 vs. ln [ATN], respectively.
Effect of temperature on rate constant was studied. Results
are shown in Table-2. Each value in the table is the average
of several runs. Experimental errors in the rate constant are
estimated to be 10 %. As expected, results show that the rate
constant (krxn2) for step 5 is greater than that for step 2 (krxn1).
The activation energy (Ea) for both steps 2 and 5, were calcu-
lated using Arrhenius equation. From plots of ln k vs. (1/T)
for both steps (Figs. 9 and 10), Ea’s of 25.64 and 20.69 KJ for
steps 2 and 5, respectively.Apparently; the thermodynamics is in
TABLE-2
REACTION RATES FOR OXIDATION OF ATROPINE BY
KMnO4 IN AQUEOUS BASIC SOLUTIONS AT VARIOUS
TEMPERATURES; [ATN] = 1.0 × 10–2 M, [MnO4–] =
1.0 × 10–3 M, [KOH] = 0.10 M
Temp. (°C)
10.0
krxn1 (dm3 mol–1 s–1)
6.21 × 10–3
krxn2 (dm3 mol–1 s–1)
2.72 × 10–2
15.0
25.0
35.0
7.85 × 10–3
3.22 × 10–2
9.59 × 10–3
3.96 × 10–2
15.90 × 10–3
5.70 × 10–2

3882 Saleh et al.
Asian J. Chem.
1/T (K–1)
3.20E-03 3.25E-03 3.30E-03 3.35E-03 3.40E-03 3.45E-03 3.50E-03 3.55E-03
-3.5
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1/T (K–1)
3.20E-03 3.25E-03 3.30E-03 3.35E-03 3.40E-03 3.45E-03 3.50E-03 3.55E-03
-2.50
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agreement and proportional to the related rate constants
determined for these steps.
Conclusion
Reaction and kinetics of the oxidation of the biologically
important atropine by alkaline KMnO4 was studied. Experi-
mental observations show that the reaction proceeds via two
stages to give the final products: tropine and phenylmalonic
acid. Two rate determining steps with different rates were
proposed, which were in agreement with obtained kinetic
result. A proposed mechanism for the process is presented.
The mechanism suggests the formation of two intermediates
in two stages of the reaction with different rates and activation
energies.
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ACKNOWLEDGEMENTS
The authors thank The Hashemite University, Zarqa,
Jordan for supporting this work.
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