Kinetics and mechanistic study of manganese(II)-catalyzed cerium(IV) oxidation of thiamine hydrochloride in aqueous perchloric acid medium by stopped flow technique

Article abstract of DOI:10.1007/s00706-013-1005-8

The kinetics of the manganese(II)-catalyzed oxidation of thiamine hydrochloride by cerium(IV) in aqueous perchloric acid medium at a constant ionic strength of 1.10 mol dm^-3 was studied spectrophotometrically at 15, 25, 35, and 45 C by the stopped flow technique. The reaction between thiamine hydrochloride and cerium(IV) in the acid medium exhibits 1:3 stoichiometry. The main products were identified by spot test, IR, ¹H NMR, and GC-MS studies. The reaction is first order in cerium(IV) and manganese(II) and has less than unit order in thiamine hydrochloride. As the acid concentration increases the rate of reaction decreases. The added product cerium(III) retards the rate of reaction. The active catalyst and oxidant species were identified as [Mn(H₂O)₄]²⁺ and [Ce(OH)]³⁺, respectively. A probable mechanism involving free radicals and the formation of a complex between substrate and catalyst is proposed. The reaction constants, activation parameters, and thermodynamic quantities are calculated and discussed.

Full text of DOI:10.1007/s00706-013-1005-8

Monatsh Chem (2013) 144:1307–1317
DOI 10.1007/s00706-013-1005-8
ORIGINAL PAPER
Kinetics and mechanistic study of manganese(II)-catalyzed
cerium(IV) oxidation of thiamine hydrochloride in aqueous
perchloric acid medium by stopped flow technique
•
•
•
D. V. Naik K. S. Byadagi S. T. Nandibewoor
S. A. Chimatadar
Received: 2 February 2012 / Accepted: 26 April 2013 / Published online: 5 June 2013
Ó Springer-Verlag Wien 2013
Abstract The kinetics of the manganese(II)-catalyzed
oxidation of thiamine hydrochloride by cerium(IV) in
aqueous perchloric acid medium at a constant ionic
strength of 1.10 mol dm-3 was studied spectrophotomet-
rically at 15, 25, 35, and 45 °C by the stopped flow
technique. The reaction between thiamine hydrochloride
and cerium(IV) in the acid medium exhibits 1:3 stoichi-
ometry. The main products were identified by spot test, IR,
1H NMR, and GC–MS studies. The reaction is first order in
cerium(IV) and manganese(II) and has less than unit order
in thiamine hydrochloride. As the acid concentration
increases the rate of reaction decreases. The added product
cerium(III) retards the rate of reaction. The active catalyst
and oxidant species were identified as [Mn(H2O)4]2? and
[Ce(OH)]3?, respectively. A probable mechanism involv-
ing free radicals and the formation of a complex between
substrate and catalyst is proposed. The reaction constants,
activation parameters, and thermodynamic quantities are
calculated and discussed.
of the seeds of many plants including cereal grains. Thiamine
is fundamentally associated with carbohydrate metabolism.
It is phosphorylated in the body to the active coenzyme thi-
amine pyrophosphate that functions as co-carboxylase for
various reactions in carbohydrate metabolism [1]. Thiamine
might also serve as a modulator of neuromuscular trans-
mission. It binds to isolated nicotinic cholinergic receptors,
and neurotransmission is impaired by pyrithiamine, a thia-
mine antimetabolite [2]. In thiamine deficiency, the
oxidation of a-ketoacid is impaired, resulting in an increase
in the concentration of pyruvate in the blood. The require-
ment for thiamine is related to the metabolic rate and is
greatest when carbohydrate is a source of energy [3]. The
oxidation kinetics and mechanism of vitamin B1 are thus
important to the process in vitro.
Manganese(II) is important in both animal and plant
enzymes and is necessary for photosynthesis. Manga-
nese(II) most probably acts as either catalyst or reductant
(the reduction potential [4] of the couple Mn(III)/Mn(II) is
1.51 V in dilute acid). Most of the studies employed
manganese(II) as a catalyst [5].
Keywords Thiamine hydrochloride ꢀ Cerium(IV) ꢀ
Manganese(II) ꢀ Oxidation ꢀ Catalysis ꢀ Kinetics
Cerium(IV) is a well-known oxidant in acid media [6]
(the reduction potential [4] of the cerium(IV)/cerium(III)
couple is 1.70 V) and is stable only in high acid concen-
trations. In sulfuric acid and sulfate media cerium(IV)
forms several sulfate complexes [7]. Similarly in perchloric
Introduction
Thiamine hydrochloride (TH) is known as vitamin B1. It was
the first member of the vitamin B complex to be isolated and
identified as a vitamin. Vitamin B1 occurs in the outer coats
acid solutions [8, 9] cerium(IV) exists as Ce4?
,
[Ce(OH)]3?, [Ce(OH2)]2?, [Ce–O–Ce]6?, and [HOCe–O–
CeOH]4?. A potentiometric study [10] of cerium(IV)
hydrolysis indicates [Ce(OH)]3? to be the predominant
active species of cerium(IV), in the concentration range
0.30–2.0 mol dm-3 HClO4, but its role has not received
much attention so far. The oxidation of organic compounds
by cerium(IV), in general, proceeds via the formation of an
intermediate complex [11]. Among the inorganic substrates
D. V. Naik ꢀ K. S. Byadagi ꢀ S. T. Nandibewoor ꢀ
S. A. Chimatadar (&)
P.G. Department of Studies in Chemistry, Karnatak University,
Pavate Nagar, Dharwad 580003, Karnataka, India
e-mail: schimatadar@gmail.com
123

1308
D. V. Naik et al.
which serve as ligands for cerium(IV) are chloride, bro-
mide, and hypophosphite [12].
decrease with increasing perchloric acid concentration
(Table 1). The order with respect to perchloric acid con-
centration was determined from the log–log plots of (rate)c
vs. acid concentrations and was found to be negative
fractional (-0.40). Cerium(IV) is known to form several
complexes in aqueous perchloric acid media [8, 9], such as
The oxidation of TH has been performed with ferri-
cyanide [13], hypoiodide [14], chlorite [15], and
N-chlorobenzenesulfonamide [16]. However, there was no
report on the oxidation of TH by cerium(IV) in the litera-
ture. Thus we tried to study this reaction in sulfuric acid
and perchloric acid media. The reaction is facile only in the
presence of micro-amounts (10-6 mol dm-3) of manga-
nese(II) in perchloric acid medium, but not in sulfuric acid.
The mechanism may be quite complicated owing to the
formation of different cerium(IV) and manganese(II)
complexes in the form of active species in the perchloric
acid medium. Hence, we have investigated the manga-
nese(II)-catalyzed oxidation of TH by cerium(IV) in order
to understand the behavior of the active species of the
oxidant, [Ce(OH)]3?, and catalyst, [Mn(H2O)4]2?, and to
propose a suitable mechanism.
Ce4?
, , , , and
[Ce(OH)]3? [Ce(OH2)]2? [Ce–O–Ce]6?
[HOCe–O–CeOH]4?. A potentiometric study of the
hydrolysis of cerium(IV) indicates [10] [Ce(OH)]3? to be
the predominant active species of cerium(IV) in perchloric
acid in the concentration range 0.30–2.0 mol dm-3
.
Effect of [manganese(II)]
At constant oxidant, reductant, and acid concentration of
2.0 9 10-4, 5.0 9 10-2, and 1.0 mol dm-3, respectively,
and I = 1.10 mol dm-3, the catalyst manganese(II) con-
centration was varied between 1.0 9 10-6 and
10.0 9 10-6 mol dm-3 (Table 1). The order with respect
to manganese(II) concentration was found to be unity. As
catalyst concentration increases the rate of the reaction also
increases, and the activity of catalyst reaches a maximum
when [manganese(II)] is at least 10.0 9 10-6 mol dm-3
under the aforementioned experimental conditions.
Results and discussion
Reaction orders
The reaction orders were determined from the slopes of
log(rate)c vs. log(concentration) plots by varying the con-
centrations of cerium(IV), TH, perchloric acid, and
manganese(II) catalyst, while keeping all other concentra-
tions and conditions constant.
Effect of ionic strength and dielectric constant
At constant concentration of reactants and other conditions
constant, the ionic strength was varied from 1.10 to
3.40 mol dm-3 by varying the concentrations of sodium
perchlorate and the rate was found to increase with
increasing ionic strength. A plot of log(rate)c vs. HI is linear
with a positive slope (Fig. 1). The effect of the dielectric
constant was studied by varying the volume fraction acetic
acid–water (u, see [17]) in the reaction mixture with all other
conditions being kept constant. It was found that (rate)c
increased with decreasing dielectric constant of the medium.
In other words, as the acetic acid content in the reaction
mixture increases, the pH of the reaction medium decreases
and the rate of the reaction [(rate)c] increases. No reaction of
the solvent occurred during the oxidant under the experi-
mental conditions employed. A plot of log(rate)c vs. 1/D is
linear with positive slope (Fig. 1).
Effect of [oxidant] and [substrate]
At constant concentration of TH (5.0 9 10-2 mol dm-3),
perchloric acid (1.0 mol dm-3), catalyst manganese(II)
(5.0 9 10-6 mol dm-3), and at constant ionic strength
(I = 1.10 mol dm-3), the cerium(IV) concentration was
varied from 0.50 9 10-4 to 5.0 9 10-4 mol dm-3; the
initial rate increased with increase in the cerium(IV) con-
centration (Table 1). From the plot of log(rate)c vs.
log[Ce(IV)], the order with respect to cerium(IV) concen-
tration was found to be unity.
The substrate TH concentration was varied in the con-
centration range 5.0 9 10-3–5.0 9 10-2 mol dm-3 at 25°
C, keeping all other conditions constant. The initial rate
increased with increase in the concentration of TH
(Table 1). From the slope of the plot of log(rate)c vs.
log[TH], the order with respect to TH concentration was
found to be fractional (0.59).
Effect of initially added products
The effect of initially added product cerium(III) was studied
in the 5.0 9 10-5–5.0 9 10-4 mol dm-3 concentration
range, keeping the ionic strength, reactant concentration, and
other conditions constant. It was found that cerium(III) in the
concentration range 5.0 9 10-5–5.0 9 10-4 mol dm-3
caused a decrease in rate. The inhibitory effect of initially
added product cerium(III) is shown in Fig. 2.
Effect of [acid]
At constant ionic strength (I = 1.10 mol dm-3) and with
other conditions remaining constant, the rate was found to
123

Kinetics and mechanistic study of manganese(II)-catalyzed cerium(IV) oxidation
1309
Table 1 Effect of variation of cerium(IV), thiamine hydrochloride, perchloric acid, and manganese(II) concentrations on the manganese(II)-
catalyzed cerium(IV) oxidation of thiamine hydrochloride in aqueous perchloric acid medium at 25 °C and I = 1.10 mol dm-3
[Ce(IV)] 9 104/mol dm-3
[TH] 9 102/mol dm-3
[HClO4]/mol dm-3
[Mn(II)] 9 106/mol dm-3
(rate)c 9 107/mol dm-3 s-1
Found
Calculated
0.50
1.0
2.0
3.0
4.0
5.0
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
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
5.0
5.0
5.0
5.0
5.0
5.0
0.5
1.0
2.0
3.0
4.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
0.1
0.2
0.4
0.6
0.8
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
0.0
1.0
2.0
3.0
4.0
5.0
6.0
7.0
8.0
9.0
10.0
1.03
2.00
3.89
5.48
7.94
9.44
1.04
1.28
2.19
2.92
3.24
3.89
9.89
7.41
5.62
4.73
4.32
3.89
0.99
2.01
3.99
5.41
7.88
9.76
1.00
1.29
2.25
2.97
3.35
3.99
9.81
7.37
5.67
4.53
4.12
3.99
Immeasurably slow
0.80
1.60
2.34
3.16
3.89
4.68
5.50
6.17
7.00
7.76
0.79
1.59
2.39
3.19
3.99
4.79
5.59
6.39
7.19
7.99
Test for free radicals
rate constant k (the slow step), K1 (the equilibrium constant
of the first equilibrium step), and K2 (the equilibrium
constant of the second equilibrium step) were obtained
from the plots of [Ce(IV)][Mn(II)]/(rate)c vs. 1/[TH] and
[Ce(IV)][Mn(II)]/(rate)c vs. [H?] (Scheme 1; Fig. 3) at
four different temperatures (Table 2). The energy of acti-
vation Ea was evaluated from the slope of the Arrhenius
plot of log k vs. 1/T (Table 2). The enthalpy of activation
DàH and the entropy of activation DàS were obtained by
using the Eyring equation [18].
The intervention of free radicals was examined as follows:
The reaction mixture, to which a known quantity of acrylo-
nitrile (scavenger) had been added initially, was kept in an
inert atmosphere for 2 h at room temperature. When the
reaction mixture was diluted with methanol, formation of
precipitate resulted, suggesting that free radicals are taking
part in the reaction.
Effect of temperature
ꢀ
ꢁ
ꢀ
ꢁ
ꢁDzG=RT
DzHþTDzS
ꢁ
kBT
h
kBT
h
The rate of reaction was measured at different temperatures
(15, 25, 35, and 45 °C) under varying acid and TH con-
centrations by using the stopped flow technique. The initial
rate was found to increase with increasing temperature. The
k ¼
e
¼
e
=RT
ð1Þ
where k is the rate constant, kB is the Boltzmann’s constant,
R is the gas constant, h is Planck’s constant, T is the
123

1310
D. V. Naik et al.
54.29 0.03 kJ mol-1. The van’t Hoff plot (log K1 vs.
1/T) was drawn and the thermodynamic quantities were
calculated (Table 2).
1/D 102
0.5
0.9
1.3
1.7
2.1
2.5
2.9
1.0
0.9
0.8
0.7
0.6
0.5
1.0
0.8
0.6
0.4
0.2
The oxidation of TH by cerium(IV) does not occur to any
appreciable extent in the aqueous sulfuric acid and perchloric
acid media under the present experimental conditions. The
reaction is facile only in the presence of micro-amounts
(10-6 mol dm-3) of manganese(II) in aqueous perchloric
acid medium, but not in sulfuric acid. In the presence of sul-
furic acid the manganese(II) catalyst is inefficient, perhaps
owing to the fact that the active species of manganese(II),
[Mn(H2O)4]2?, does not form in sulfuric acid media. Hence
the study was undertaken in aqueous perchloric acid medium.
In perchloric acid medium the reaction between TH and cer-
ium(IV) in the presence of manganese(II) has a stoichiometry
of 1:3, i.e., 1 mole of TH reacts with 3 moles of cerium(IV) to
give the final products cerium(III), 4-amino-2-methylpyrimi-
dine-5-carbaldehyde, and 2-(4-methylthiazol-5-yl)ethanol.
The reaction is first order with respect to cerium(IV) and
manganese(II) concentrations and less than unit order (0.59)
with respect toTH. Theeffectofhydrogenionsonthe ratewas
studied by adding perchloric acid and it was found that as the
perchloric acid concentration increased in the reaction mix-
ture the rate of the reaction decreased. The order with respect
to perchloric acid concentration is less than unity and nega-
tive. Similar behavior was reported for the oxidation of
mandelic acid [19] and mercury(I) [20] by cerium(IV) in
perchloric acid medium. Initially added product cerium(III)
retards the reaction rate appreciably. Similar results were
obtained in the oxidation of tellurium(IV) [21], L-glutamic
acid [22], and glycerol [23] by cerium(IV). The experimental
result suggest that TH combines with the active species of the
catalyst to form a complex, which then reacts in a slow step
with 1 mole of cerium(IV) to give the product cerium(III), a
free radical derived from TH, and an unstable complex, with
generation of the catalyst manganese(II). The free radical
derived from TH undergoes hydrolysis to give another inter-
mediate product, 4-amino-2-methylpyrimidine-5-carbinol.
This intermediate reacts with 2 moles of cerium(IV) in a
further fast step to give the final product cerium(III) and
4-amino-2-methylpyrimidine-5-carbaldehyde. The interme-
diate complex formed in the rate-determining step combines
with H? ions to give the other final product 2-(4-methylthia-
zol-5-yl)ethanol in further fast step. The experimental results
are accommodated in the form of the following mechanism
(Scheme 1).
0.5
0.8
1.1
1.4
1.7
2.0
I
Fig. 1 Effect of ionic strength and dielectric constant on the
manganese(II)-catalyzed cerium(IV) oxidation of thiamine hydro-
chloride at 25 °C
Fig. 2 Effect of added product cerium(III) on the manganese(II)-
catalyzed oxidation of thiamine hydrochloride by cerium(IV) at 25 °C
absolute temperature, and DàG is the free energy of
activation. The linear form of Eq. (1) is
z
RT
z
k
D H D S
kB
h
ln ¼ ꢁ
þ
þ ln
R
ð2Þ
T
The results indicate the formation of a complex between
TH and manganese(II). The spectral evidence for the for-
mation of a complex between the substrate and catalyst was
obtained from UV–Vis spectra of TH and thiamine hydr-
chloride–manganese(II) mixtures in which a hypsochromic
shift of 4 nm from 284 to 280 nm was observed. The
formation of a complex was also proved kinetically by a
The slope of the plot of log k/T vs. 1/T gives the value of
enthalpy of activation DàH (Table 2). By using the value of
DàH and the rate constant at a particular temperature T the
value of DàS was obtained by simple rearrangement of
Eq. (1) (Table 2). Using these values of DàH and DàS, the
free energy of activation DàG298 is obtained as
123

Kinetics and mechanistic study of manganese(II)-catalyzed cerium(IV) oxidation
1311
Scheme 1
K1
H+
Ce4+
.
Ce4+ H+
+
_
Cl
+
N
CH3
CH2
K2
N
Complex (C1)
+
Mn(II)
S
CH2CH2OH
NH2
H3C
N
.
CH2 Cl
+
CH3
k
N
N
Ce4+
Ce3+
+
+
Complex (C1)
Mn(II)
+
+
slow
H3C
N
NH2
CH2CH2OH
S
.
CH2OH
CH2
N
N
fast
H+
N
+
+
HOH
H3C
NH2
H3C
N
NH2
CHO
CH2OH
NH2
N
fast
N
2 Ce4+
2 H+
+
2 Ce3+
+
+
H3C
N
NH2
H3C
N
+
Cl
CH3
CH3
CH2CH2OH
fast
N
N
H+
CH2CH2OH
+
HCl
+
S
S
Therefore,
Michaelis–Menten plot (Fig. 3). Such formation of a
complex between substrate and catalyst was also observed
in literature [22, 23]. Scheme 1 leads to the rate law Eq. (8)
which explains the observed order with respect to cer-
ium(IV), TH, perchloric acid, and manganese(II):
½THꢂt
½THꢂf ¼
ð4Þ
1 þ K2½MnðIIÞꢂ
where subscripts t and f stand for total and free,
respectively. In the experiment a very low concentration
of manganese(II) is used. Hence, the term K2[Mn(II)] may
be neglected in comparison with unity in Eq. (4). So,
ðrateÞc ¼ k½complexðC1Þꢂ ½CeðIVÞꢂ
ð3Þ
kk1k2½CeðIVÞꢂ½THꢂ½MnðIIÞꢂ
¼
½Hþꢂ
½THꢂf ¼ ½THꢂt
ð5Þ
Similarly,
The total concentration of thiamine hydrochloride [TH]t
is given by
½MnðIIÞꢂt ¼ ½Mnð IIÞꢂ þ K2½THꢂ ½MnðIIÞꢂ
¼ ½MnðIIÞꢂ ff 1 þ K2½THꢂg
½THꢂt ¼ ½THꢂ þ ½C1ꢂ
¼ ½THꢂf þ K2½THꢂ ½MnðIIÞꢂ
f
ð6Þ
½MnðIIÞꢂt
½MnðIIÞꢂt ¼
f
¼ ½THꢂff1 þ K2½MnðIIÞꢂg
1 þ K2½THꢂ
f
123

1312
D. V. Naik et al.
Table 2 Effect of temperature on the manganese(II)-catalyzed cer-
ium(IV) oxidation of thiamine hydrochloride in aqueous perchloric
acid medium
Temperature/°C
k 9 10-2/s-1
Rate constant with respect to slow step of Scheme 1
15
25
35
45
15.54
19.44
23.12
27.32
Parameters
Value
Activation parameters
DàH/kJ mol-1
DàS/J K-1 mol-1
DàG/kJ mol-1
log A
11.77 0.46
-142.6 1.5
54.29 0.03
5.7 0.05
Temperature/°C
K1/dm3 mol-1
K2 9 105/dm3 mol-1
Fig. 3 Verification of rate law Eq. (8) in the form of Eq. (9)
Effect of temperature on first and second equilibrium steps of
Scheme 1
and
½CeðIVÞꢂt ¼ ½CeðIVÞꢂf þ ½CeðIVÞꢂ½Hþꢂ
15
25
35
45
0.82
0.75
0.56
0.41
26.21
18.46
12.43
8.10
½CeðIVÞꢂf
¼ ½CeðIVÞꢂf þ K1
½Hþꢂ
ꢂ
ꢃ
K1
½Hþꢂ
¼ ½CeðIVÞꢂf 1þ
Values from K1
Values from K2
Thermodynamic quantities with respect to first and second
equilibrium steps of Scheme 1
Therefore,
½CeðIVÞꢂt
DH/kJ mol-1
DS/J K-1 mol-1
DG/kJ mol-1
-18 3.2
-63.4 10.6
0.94 0.23
-29.8 1.6
-172
21.5 0.1
½CeðIVÞꢂf ¼
ð7Þ
K1
1 þ
Hþ
5
Substituting Eqs. (5, 6, 7) in Eq. (3) and omitting the
subscripts, we have
½Hþꢂ
1
ꢁd½CeðIVÞꢂ
ðInterceptÞ1 ¼
þ
ð11Þ
ð12Þ
ðrateÞc¼
k K1
k
dt
k K1 K2 ½CeðIVÞꢂ ½THꢂ ½MnðIIÞꢂ
Therefore,
¼
ð8Þ
½Hþꢂ þ K1 K2 ½THꢂ ½Hþꢂ þ K1 K2 ½THꢂ
ðInterceptÞ1
¼ K2
¼ Formation constant of the complexðC1Þ
ðSlopeÞ1
The rate law Eq. (8) may be rearranged to Eq. (9) which
is suitable for verification.
½CeðIVÞꢂ ½MnðIIÞꢂ
ðrateÞc
½Hþꢂ
1
½Hþꢂ
1
II. From the second plot, i.e., [Ce(IV)] [Mn(II)]/(rate)c
vs. 1/[H?], we obtain
¼
þ
þ
þ
k K1 K2½THꢂ k K2½THꢂ k K1
k
ð9Þ
1
1
ðSlopeÞ2 ¼
þ
k K1 K2½THꢂ
k K1
ꢃ
According to Eq. (9) the plots of [Ce(IV)] [Mn(II)]/
(rate)c vs. 1/[TH] and [Ce(IV)] [Mn(II)]/(rate)c vs. [H?]
should be linear and are found to be so (Fig. 3).
I. From the first plot, i.e., [Ce(IV)] [Mn(II)]/(rate)c vs.
1/[TH], we obtain
ꢂ
1
K2½THꢂ þ 1
¼
ð13Þ
ð14Þ
k K1
K2½THꢂ
1
1
ðInterceptÞ2 ¼
þ
k K2½THꢂ
k
ꢂ
ꢃ
½Hþꢂ
1
1
1 þ K2½THꢂ
K2½THꢂ
ðSlopeÞ1¼
þ
ð10Þ
¼
k K1 K2 k K2
k
123

Kinetics and mechanistic study of manganese(II)-catalyzed cerium(IV) oxidation
1313
Hence,
Experimental and calculated values agree reasonably well
(Table 1), which fortifies the proposed scheme.
ðInterceptÞ2
¼ K1
The active species involved in the mechanism can be
understood as follows: cerium(IV) is known to form sev-
eral species in aqueous perchloric acid media [8, 9], such
as Ce4?, [Ce(OH)]3?, [Ce(OH2)]2?, [Ce–O–Ce]6?, and
[HOCe–O–CeOH]4?.The potentiometric study of the
hydrolysis of cerium(IV) indicates that [Ce(OH)]3? is the
predominant active species of cerium(IV) in perchloric
acid [10]. Hence, [Ce(OH)]3? is considered as the active
oxidant species. The manganese(II) catalyst is understood
to be present in the aqueous acid medium mainly as
[Mn(H2O)4]2? [24]. Manganese(II)-catalyzed oxidation of
TH is greatly facilitated in perchloric acid solution and the
reason may be that cerium(IV) species such as [Ce(OH)]3?
and manganese(II) [Mn(H2O)4]2? are actively involved.
Slope2
¼ Ionization constant of Ce4þ ꢀ Hþ
ð15Þ
Substituting the values of K1 and [H?] in the (intercept)1
of the first plot, i.e., in Eq. (11), we get the value of k, the
rate constant with respect to the slow step of Scheme 1. By
this procedure the obtained values of k, K1, and K2 are
1,944.4 50 mol dm-3 s-1, 0.75 0.04 dm3 mol-1, and
18.46 0.40 dm3 mol-1, respectively, at 25 °C. The
value of K1 obtained is in reasonable agreement with the
literature [20] (K1 = 0.56). Using these values, we
calculated rate constants (rate)c under different
experimental conditions by applying the rate law Eq. (8)
and compared the results with the experimental data.
K3
Scheme 2
.
.
Ce4+
+
H2O
Ce(OH)3+ H+
K4
[ Mn(H2O)4 ] 2+
H O
4
MnSO4
+
2
K5
.
Ce(OH)3+
H+
Ce(OH)3+ H+
+
_
Cl
S
+
N
CH3
CH2
NH2
K6
N
+
[ Mn(H2O)4 ]2+
Complex (C2)
+
CH2CH2OH
H3C
N
.
CH2 Cl
+
CH3
k
N
N
Ce(OH)3+
Ce(OH)2+
H3C
+
[Mn(H2O)4]2+
Complex (C2)
+
+
slow
N
NH2
CH2CH2OH
S
.
CH2OH
CH2
N
fast
H+
N
+
+
HOH
H3C
N
NH2
H3C
N
N
NH2
CHO
CH2OH
NH2
N
fast
+
+
2 H+
2 Ce(OH)3+
Ce(OH)2+
+
2
H3C
N
NH2
H3C
N
+
Cl
CH3
CH3
CH2CH2OH
fast
N
N
H+
CH2CH2OH
+
HCl
+
S
S
123

1314
D. V. Naik et al.
The mechanism in Scheme 1 will therefore involve the
active species as shown in Scheme 2.
Experimental
Therefore, in terms of active species, the rate law
Eq. (8) takes the following form (Eq. (16)):
The solutions were prepared in doubly distilled water and
reagent grade chemicals were used. The stock solution of
the oxidant cerium(IV) was obtained by dissolving cer-
ium(IV) ammonium sulfate (E. Merck) in 1.0 mol dm-3 of
H2SO4 and standardized with iron(II) ammonium sulfate
solution [28]. A stock solution of TH (HIMEDIA) was
prepared by dissolving TH in distilled water. Manga-
nese(II) stock solution was made by dissolving a known
weight of MnSO4 (sd fine-chem Ltd.) in distilled water.
Acetic acid and perchloric acid were obtained from Glaxo
Excelar. The former was purified by refluxing with potas-
sium permanganate for 5–6 h and then distilling; a further
distillation yielded the fraction of b.p. 118 °C for use.
Cerium(III) solution was prepared by dissolving cer-
ium(III) sulfate (BDH) in water. Sodium perchlorate and
perchloric acid were used to provide the required ionic
strength and acidity, respectively. Sodium perchlorate was
prepared by mixing equal volumes of sodium hydroxide
and perchloric acid and neutralizing to pH 7.0.
2þ
kK5K6½CeðOHÞ3þꢂ½THꢂ½MnðH2O4Þꢂ
ðrateÞc ¼
½HþꢂK5 þ K6½THꢂ½Hþꢂ þ K5K6½THꢂ
kK3K4K5K6½CeðOHÞ3þꢂ½THꢂ½MnðH2O4Þꢂ
2þ
¼
2ꢁ
½Hþꢂ½SO4ꢂ f½Hþꢂ þ K5 þ K6½THꢂ½Hþꢂ þ K5K6½THꢂg
ð16Þ
The effect of ionic strength on the rate qualitatively
accounts for the reaction between two positively charged
ions as seen in Scheme 1. The effect of solvent on the
reaction rate has been described in detail in the literature
[25]. The increase in the acetic acid content in the reaction
medium leads to an increase in the rate of reaction which
might be due to the presence of ionic species of TH and the
dissociated form of the acetate ion. In the present study, the
plot of log(rate)c vs. 1/D is linear with a positive slope,
which is in the right direction as shown in Scheme 1.
The values of DàH and DàS are both favorable for
electron transfer processes. The negative values of
DàS indicate that the complex is more ordered than the
reactants [5]. The value of DàS is within the range of those
expected for radical reactions and has been ascribed to the
nature of the electron pairing and unpairing process and to
the loss of degree of freedom formerly available to the
reactants upon formation of a rigid transition state [26].
The observed modest enthalpy of activation as well as the
high value of the rate constant of the slow step indicates
that the oxidation presumably occurs via an inner-sphere
mechanism. This conclusion is supported by the literature
[27].
For kinetic measurements a UV–Vis spectrophotometer
(Varian CARY 50 Bio) connected to a rapid kinetic
accessory (HI-TECH SFA-12) was used. For product
analysis, GC–MS data was obtained on a Shimadzu 17A
gas chromatograph with a Shimadzu QP-5050A mass
spectrometer using EI ionization, IR spectral data was
obtained on a Nicolet 5700-FT-IR spectrometer (Thermo,
1
USA), H NMR data was obtained on a Bruker 300 MHz
spectrometer, and for pH measurements an Elico pH meter
model LI120 was used.
2.5
2.0
Conclusion
The reaction is immeasurably slow in perchloric acid at
room temperature. However, the reaction is facile in the
presence of micro-amounts (106 mol dm-3) of manga-
nese(II). Manganese(II)-catalyzed cerium(IV) oxidation of
thiamine hydrochloride (TH) in perchloric acid has a
reductant to oxidant stoichiometry of 1:3. The order with
respect to both oxidant and catalyst is found to be unity,
whereas the order with respect to substrate is less than
unity. The products were identified by UV–Vis, GC–MS,
1.5
1.0
0.5
0.0
A
B
C
1
IR, and H NMR spectra. The added product cerium(III)
retards the rate of reaction. The active species of oxidant
and catalyst are [Ce(OH)]3? and [Mn(H2O)4]2?, respec-
tively. The decomposition of the complex is the slow step
in the reaction mechanism and is followed by fast steps to
give the products.
0
5
10
15
20
25
Time/s
Fig. 4 Plot of cerium(IV) concentration vs. time at 25 °C
123

Kinetics and mechanistic study of manganese(II)-catalyzed cerium(IV) oxidation
1315
Kinetic measurements
perchloric acid, sodium perchlorate, and manganese(II)
catalyst because the initial reaction was too fast to be mon-
itored by usual methods. The course of the reaction was
followed by monitoring the decrease in absorbance of cer-
ium(IV) in a 1-cm quartz cell placed in the thermostated
compartment of a the UV–Vis spectrophotometer at its
absorption maximum of 360 nm as a function of time. The
All kinetic runs were carried out under pseudo-first-order
conditions where TH concentration was in excess over cer-
ium(IV) concentration at a constant temperature of
25 0.1 °C. The thermally equilibrated solution of cer-
ium(IV) and TH also contained the required quantities of
Scheme 3
_
Cl
+
CH3
CH2
NH2
Mn(II)
N
N
3 Ce4+
+
H2O
N
+
HClO4
S
CH2CH2OH
H3C
N
CHO
NH2
CH3
CH2CH2OH
N
3 Ce3+
+
+
3H+
+
+
HCl
S
H3C
N
Fig. 5 IR spectra of the product 4-amino-2-methylpyrimidine-5-carbaldehyde
123

1316
D. V. Naik et al.
perchloric acid and sodium perchlorate provide the required
acidity and ionic strength in the reaction medium. The
application of Beer’s law under the reaction conditions had
been verified at 360 nm in 1.0 mol dm-3 perchloric acid.
The molar absorptivity index of cerium(IV) at 360 nm was
found to be e = 3,500 50 dm3 mol-1 cm-1. In the pres-
ent study one of the products cerium(III) retards the reaction;
hence, initial rates were determined from a plot of [cer-
ium(IV)] vs. time (Fig. 4) by employing the plane mirror
method [29, 30]. The initial rates (rate)c obtained were
reproducible within 5 % and are the average of at least four
independent kinetic runs.
for 8 h at room temperature in an inert atmosphere. The
unreacted cerium(IV) concentration was assayed by mea-
suring the absorbance at 360 nm. The products were
1
identified and confirmed by UV–Vis, IR, GC–MS, and H
NMR spectra as cerium(III), 4-amino-2-methylpyrimidine-
5-carbaldehyde, and 2-(4-methylthiazol-5-yl)ethanol. The
results showed that 2-methylpyrimidine-5-carbaldehyde
was a major product in the reaction. The results indicate
that 1 mole of TH reacted with 3 moles of cerium(IV)
according to Scheme 3.
The product cerium(III) was confirmed by UV–Vis
spectra at 230–260 nm. The products 4-amino-2-methyl-
pyrimidine-5-carbaldehyde and 2-(4-methylthiazol-5yl)
ethanol were identified by spot tests [31]. 4-Amino-2-
methylpyrimidine-5-carbaldehyde was identified by IR
spectra owing to the presence of the C=O stretch at
1,662 cm-1 and –NH2 stretch at 3,432 cm-1 (Fig. 5).
4-Amino-2-methylpyrimidine-5-carbaldehyde and 2-(4-
methylthiazol-5-yl)ethanol were confirmed by GC–MS
Stoichiometry and product analysis
Different reaction mixtures with different sets of concen-
trations of reactants, where cerium(IV) concentration was
in excess over TH concentration, at constant ionic strength,
acidity, and at constant concentrations of catalyst were kept
Fig. 6 GC–MS spectra of the product 4-amino-2-methylpyrimidine-5-carbaldehyde
Fig. 7 GC–MS spectra of the
product 2-(4-methylthiazlole-5-
yl)ethanol
123

Kinetics and mechanistic study of manganese(II)-catalyzed cerium(IV) oxidation
1317
11. Guilbault GG, McCurdy WH Jr (1963) J Phys Chem 67:283
spectra owing to the presence of molecular ion peaks at
134 amu (Fig. 6) and 143 amu (Fig. 7), respectively. All
other peaks in the GC–MS specta were consistent with the
observed structures. 4-Amino-2-methylpyrimidine-5-car-
12. Carroll RL, Thomas LB (1966) J Am Chem Soc 88:1376
13. Garger G, Barger F, Todd AR (1935) Nature 136:259
14. Jaronsanti J, Panijan B (1981) Experientia 37:248
15. Kawasaki C, Horio TJ (1954) Pharm Soc Jpn 74:904
16. Mohana KN, Prasad N, Rai KML (2008) Monatsh Chem
139:1203
1
baldehyde was also confirmed by H NMR spectra, which
were recorded on a Bruker 300 MHz spectrometer using
DMSO-d6 as solvent and TMS as internal reference. The
product showed a singlet at d = 2.34 ppm accounting for
the three protons of the methyl group at position 2. The
amine group appeared as a broad singlet (D2O exchanged)
at 5.33 ppm. The proton at the 6-position and the aldehyde
proton appeared as singlets at 8.06 and 9.46 ppm,
respectively.
¨
17. Gamsjager H, Lorimer JW, Scharlin P, Shaw DG (2008) Pure
Appl Chem 80:233
18. Byadagi KS, Naik DV, Savanur AP, Nandibewoor ST, Chi-
matadar SA (2010) Reac Kinet Mech Cat 99:53
19. Sengupta KK, Aditya S, Ghosh BN (1963) J Indian Chem Soc
40:1823
20. McCurdy WH Jr, Guilbault GG (1960) J Phys Chem 64:1825
21. Dikshitulu LSA, Hanumantarao DSN (1981) Indian J Chem
20A:784
22. Chimatadar SA, Madawale SV, Nandibewoor ST (2007) Transit
Met Chem 32:634
23. Savanur AP, Nandibewoor ST, Chimatadar SA (2009) Transit
Met Chem 34:711
References
24. Amis ES (1996) Solvent effects on reaction rates and mechanism.
Academic, New York
25. Entelis SG, Tiger RP (1976) Reaction kinetics in the liquid phase.
Wiley, New York
26. Walling C (1976) Free radicals in solutions. Academic, New
York, p 38
27. Farokhi SA, Nandibewoor ST (2003) Tetrahedron 59:7595
28. Jeffery GH, Bassett J, Mendham J (2002) Vogel’s textbook of
quantitative chemical analysis, 6th edn. Longman, Essex, p 426
29. Laidler KJ (1975) Chemical kinetics, 2nd edn. McGraw-Hill,
New Delhi, p 16
30. Casado J, Lopez-Quintella MA, Lorenzo-Barral FM (1986) J
Chem Educ 63:450
31. Feigl F (1975) Spot tests in organic analysis. Elsevier, New York,
p 435
1. Dougherty HW, Pollak PI (1970) Medicinal chemistry, 3rd edn.
Wiley-Interscience, New York 30
2. Waldenlind I (1978) Acta Physiol Scand 1:459
3. Marcus R, Coulsfon AM (1991) The pharmacological basis of
therapeutics, 8th edn. Pergamon press, New York, Chap 63
4. Day MC, Selbin J (1964) Theoretical inorganic chemistry.
Reinhold, New York, p 226
5. Patil RK, Chimatadar SA, Nandibewoor ST (2008) Transit Met
Chem 33:625
6. Thabaj KA, Chimatadar SA, Nandibewoor ST (2007) Transit Met
Chem 32:902
7. Kharzeoua SE, Senebrennikou VV (1967) Russ J Inorg Chem
12:1601
8. Hardwick TJ, Robertson E (1951) Can J Chem 29:818
9. Datt N, Nagori RR, Meharotra RN (1986) Can J Chem 64:19
10. Baker FB, Newton TW, Khan M (1960) J Phys Chem 64:109
123