Kinetics and mechanism of ruthenium(III) catalyzed oxidation of chloromphenicol - An antibiotic drug by diperiodatocuprate(III) in aqueous alkaline medium

Article abstract of DOI:10.1134/S0023158411060085

The kinetics of ruthenium(III) catalyzed oxidation of chloramphenicol (CHP) by diperiodatocuprate(III) (DPC) in aqueous alkaline medium at a constant ionic strength of 0.1 mol l^-1 was studied spectrophotometrically. The reaction between DPC and CHP in alkaline medium exhibits 1: 2 stoichiometry (CHP: DPC). The main oxidation products were identified by spot test, IR, NMR, and GC-MS spectral studies. The reaction is first order with respect to ruthenium(III) and DPC concentrations. The order with respect to chloramphenicol concentration varies from first order to zero order as the chloramphenicol concentration increases. As the alkali concentration increases the reaction rate increases with fractional order dependence on alkali concentration. Increase in periodate concentration decreases the rate. A mechanism adequately describing the observed regularities is proposed. The reaction constants involved in the different steps of the mechanism were calculated. The activation parameters with respect to limiting step of the mechanism are computed and discussed. Thermodynamic quantities are determined. Pleiades Publishing, Ltd., 2012.

Full text of DOI:10.1134/S0023158411060085

ISSN 0023ꢀ1584, Kinetics and Catalysis, 2012, Vol. 53, No. 1, pp. 65–74. © Pleiades Publishing, Ltd., 2012.
Kinetics and Mechanism of Ruthenium(III) Catalyzed Oxidation
of Chloromphenicol—An Antibiotic Drug by Diperiodatocuprate(III)
in Aqueous Alkaline Medium1
R. V. Hosahalli, K. S. Byadagi, S. T. Nandibewoor, and S. A. Chimatadar
Post Graduate Department of Studies in Chemistry, Karnatak University, Pavate Nagar, Dharwad 580003, India
eꢀmail: schimatadar@gmail.com
Received February 23, 2011
Abstract—The kinetics of ruthenium(III) catalyzed oxidation of chloramphenicol (CHP) by diperiodatocuꢀ
prate(III) (DPC) in aqueous alkaline medium at a constant ionic strength of 0.1 mol l–1 was studied spectroꢀ
photometrically. The reaction between DPC and CHP in alkaline medium exhibits 1 : 2 stoichiometry
(CHP : DPC). The main oxidation products were identified by spot test, IR, NMR, and GCꢀMS spectral
studies. The reaction is first order with respect to ruthenium(III) and DPC concentrations. The order with
respect to chloramphenicol concentration varies from first order to zero order as the chloramphenicol conꢀ
centration increases. As the alkali concentration increases the reaction rate increases with fractional order
dependence on alkali concentration. Increase in periodate concentration decreases the rate. A mechanism
adequately describing the observed regularities is proposed. The reaction constants involved in the different
steps of the mechanism were calculated. The activation parameters with respect to limiting step of the mechꢀ
anism are computed and discussed. Thermodynamic quantities are determined.
DOI: 10.1134/S0023158411060085
1
In recent years the study of highest oxidation states between different copper(III) species are involved,
and it would be of interest to know, which of the speꢀ
cies is the active oxidant.
Ruthenium(III) is an efficient catalyst in many
redox reactions [16] due to its ability to form intermeꢀ
diate complexes, free radicals, and multiple oxidation
of transition metals has intrigued many researchers.
Transition metals in a higher oxidation state can be
stabilized by chelating with suitable polydentate
ligands. Metal chelates such as diperiodatocuꢀ
prate(III) [1], diperiodatoargentate(III) [2], diperioꢀ
datonickelate(IV) [3] are good oxidants in a medium
with an appropriate pH value. Periodate and tellurate
complexes of copper in its trivalent state have been
extensively used in the analysis of several organic comꢀ
state of ruthenium [16, 17]. Chloramphenicol (CHP
)
is effective against a wide variety of Gramꢀpositive and
Gramꢀnegative bacteria including most anaerobic
organisms. It is considered as a prototypical broadꢀband
pounds [4]. The kinetics of selfꢀdecomposition of spectrum antibiotic, alongside the tetracyclines. The
these complexes was studied in detail [5]. Copper(lII) clinical application of chloramphenicol was in the
treatment of typhoid. In view of potential pharmaceutiꢀ
cal importance of chloramphenicol and lack of literaꢀ
ture on the oxidation of this drug by any oxidant and the
complexity of the reaction, a detailed study of its oxidaꢀ
tion becomes important. Here, the results on the CHP
oxidation by DPC is presented, and monoperiodatocuꢀ
prate (MPC) is understood as an active intermediate.
is shown to be an intermediate in the copper(II)ꢀcataꢀ
lyzed oxidation of amino acids by peroxydisulphate
[6]. The oxidation reaction usually involves the copꢀ
per(II)–copper(I) couple, and such aspects are dealt
in different reviews [7, 8]. The use of diperiodatocuꢀ
prate(III) (DPC) as an oxidant in alkaline medium is
new and restricted to a few cases due to its limited solꢀ
ubility in aqueous medium. DPC is a versatile oneꢀ
electron oxidant for various organic compounds in
alkaline medium and its use as an analytical reagent is
well recognized [9]. Copper complexes have occupied
a major place in oxidation chemistry due to their abunꢀ
dance and relevance in biological chemistry [10–14].
Copper(III) is involved in many biological electron
transfer reactions [15]. When copper(III) periodate
complex is an oxidant, multiple equilibrium steps
EXPERIMENTAL
All chemicals used were of A.R. grade, and double
distilled water was used. The solution of chloramꢀ
phenicol (“Sisco Chem.”) was prepared by dissolving
known amount of the samples in distilled water. The
purity of the samples was checked by their melting
point (150
°
C
). (Literarure: 149–153 C). Rutheꢀ
°
nium(III) stock solution is made by dissolving a
known weight portion of RuCl3 (“SD FineꢀChem
1
The article is published in the original.
65

66
HOSAHALLI et al.
was initiated by adding DPC to chloramphenicol
3 + log
3.0
A
which also contained the required concentration of
ruthenium(III), KNO3, KOH and KIO4. Since the
initial reaction was too fast to be monitored by usual
methods, the course of the reaction was followed by
monitoring the decrease in the DPC absorbance at its
absorption maximum of 415 nm as a function of time
in a quartz cell with an optical path length of 1 cm
placed in a thermally controlled compartment of a
Varian Caryꢀ50 Bio UVꢀvis spectrophotometer conꢀ
nected to a rapid kinetic accessory (SFAꢀ12, “Hiꢀ
Tech”). Earlier the extinction coefficient was deterꢀ
2.5
2.5
2.5
2.5
2.5
1
2
3
4
5
6
mined at 415 nm for different concentrations ( =
ε
6235 250 l mol–1 cm–1). It was verified that there is
negligible interference from other species present in
the reaction mixture at this wavelength.
0
0.5
1.0
1.5
2.0
2.5 3.0
Time, min
The pseudoꢀfirst order rate constants, kobs, were
determined from log
absorbance at 415 nm, and
A
vs.
t
plots (Fig. 1), where
A is the
Fig. 1. First order plots of oxidation of chloramphenicol by
–4
–7
–
t
is the reaction time. The
DPC at 25
°C, [CHP] = 5.0 × 10 mol/l, [OH ]] =
0.05 mol/l, [Ru(III)] = 5.0
×
10
0.10 mol/l; [DPC] 10 mol/l: 0.5 (
), 4.0 ( ) and 5.0 (6).
mol/l, and
I =
plots were linear up to 80% completion of the reaction.
The rate constants were reproducible within 5% and
are the average of at least three independent kinetic
runs. In the kinetic experiments, a constant concenꢀ
5
×
1
), 1.0 ( ), 2.0 (3), 3.0
2
(4
5
tration of KIO4 (5.0
10–5 mol/l) was used throughout
×
Ltd.”, India) in 0.20 mol/l HCl. Mercury is added to
ruthenium(III) solution to reduce any ruthenium(IV)
formed during the preparation of ruthenium(III)
stock solution and kept for a day. The ruthenium(III)
concentration is assayed by EDTA titration [18].
Dilute solution of ruthenium(III) are made from the
stock solution as required. The copper(III) periodato
complex was prepared [19] and standardized by a stanꢀ
dard procedure [20]. The copper(II) solution was
made by dissolving the known amount of copper sulꢀ
phate (“BDH”, India) in distilled water. Periodate
solution was prepared by dissolving the required
amount of sample in hot water and used after 24 h. Its
concentration was ascertained iodometrically [21] at
neutral pH adjusted by a phosphate buffer. KOH and
KNO3 solutions were employed to maintain the
required alkalinity and ionic strength, respectively.
the study unless otherwise stated, and excess periodate
was present in DPC. The possibility of the chloramꢀ
phenicol oxidation by periodate in alkaline medium at
25 C was verified. It was found that there is no signifiꢀ
°
cant reaction between CHP and periodate under the
experimental conditions employed compared to the
DPC oxidation of chloramphenicol. The concentraꢀ
tion of periodate and OH– was calculated by considerꢀ
ing the amount present in the DPC solution and that
additionally added. Kinetic runs were also carried out
in N2 atmosphere in order to understand the effect of
dissolved oxygen on the reaction. No significant difꢀ
ference in the results was obtained in the presence and
in the absence of nitrogen.
In view of ubiquitous contamination with carbonꢀ
ate in the basic medium, the effect of carbonate was
also studied. Added carbonate has no effect on the
reaction rates. The spectral changes at 415 nm during
the reaction show (Fig. 2) that the concentration of
DPC decreases.
Instruments Used
For kinetic measurements, a Peltier Accessory
(temperature control) attached to Varian CARY 50
Bio UVꢀvisible spectrophotometer (“Varian”, Victoꢀ
riaꢀ3170, Australia) connected to a rapid kinetic
accessory SFAꢀ12 (“HiꢀTech Scientific Ltd.”, U.K.)
was used.
RESULTS AND DISCUSSION
Stoichiometry and Product Analysis
For product analysis, a chromatoꢀmass spectromeꢀ
ter QPꢀ2010S (“Shimadzu”), Nicolet 5700 FTꢀIR
Different sets of reaction mixtures containing
excess of DPC to chloramphenicol in the presence of
constant amount of OH–, ruthenium(III), KNO3,
KIO4 were kept for 6 h in a closed vessel under inert
atmosphere. The remaining DPC concentration was
1
spectrometer (“Thermo”, USA), 300 MHz H NMR
spectrophotometer (“Bruker”, Switzerland) were used.
For pH measurements, an LIꢀ120 pHꢀmeter
(“Elico Ltd.”, India) was used.
The ruthenium(III) catalyzed oxidation of CHP by estimated by spectrophotometrically at 415 nm. The
DPC was followed under pseudoꢀfirst order condiꢀ results, indicated that one mole of CHP requires two
tions at [CHP] > [DPC] at 25 0.1
°
C. The reaction moles of DPC as shown in Eq. (I).
KINETICS AND CATALYSIS Vol. 53
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KINETICS AND MECHANISM OF RUTHENIUM(III) CATALYZED OXIDATION
67
H
C
NHCOCHCl2
C CH2OH
1.0
O2N
OH H
+ 2[Cu(H2IO6)(H2O)2] + 2OH
–
1
2
(I)
H
C
O
NH2
HO C CH2OH
H
Ru(III)
+
O2N
3
4
+ Cl2CHCOOH + 2Cu2+ + 2H2IO63– + 4H2O.
5
After completion of the reaction, the reaction mixꢀ
ture was acidified, concentrated, and extracted with
ether. The ether layer was subjected to column chroꢀ
matography [22], and the fractions were subjected to
spectral investigations. From the IR, GCꢀMS, and
1H NMR spectra, the main oxidation product was
350.0
400.0
450.0
, nm
500.0
550.0
λ
Fig. 2. Spectral changes during the oxidation of chloramꢀ
5
identified as
pꢀnitrobenzaldehyde. IR spectrum
phenicol by DPC at 25
°
C, [DPC] = 5.0
×
10 mol/l,
10 mol/l, [OH ] = 0.05 mol/l, [Ru(III)] =
10 mol/l, and = 0.10 mol/l (scanning time interꢀ
–4
–
[CHP] = 5.0
×
showed C=O stretching mode for aldehyde functional
–7
group at 1708 cm–1, while –NO2 stretching mode was
5.0
×
I
val is 1.0 min).
observed at 1349 cm–1. The presence of
pꢀnitrobenzalꢀ
dehyde was also confirmed by GCꢀMS analysis
obtained on a Shimadzu’s 17A gas chromatograph
with a Shimadzu’s XPꢀ5000A mass spectrometer using
electron impact (EI) ionization technique. The mass
spectrum showed the base peak at 151 amu consistent
with the molecular ion of 151 amu. All other peaks
observed in the GCꢀMS can be interpreted in accorꢀ
2.5
2.0
1.5
1.0
0.5
dance with the structure of
ther
pꢀnitrobenzaldehyde. Furꢀ
p
ꢀnitrobenzaldehyde was characterized by its
1H NMR spectrum (DMSO): 10.17 (s, 1H, CHO),
8.18 (d, 2H, ArH), 7.67 (d, 2H, ArH). Another prodꢀ
uct 2ꢀaminoꢀ2ꢀhydroxyethanol was confirmed by GCꢀ
MS spectrum, which showed the molecular ion peak
at 77 amu. Dichloroacetic acid was confirmed by spot
test [23].
0
5
10
[CHP]
15
20
25
30
×
104, mol/l
Reaction Orders
The reaction orders were determined from the
slopes of the kobs vs. concentration plots in logarithmic
coordinates by varying the concentrations of chloramꢀ
Fig. 3. Dependence of
k
on [CHP].
obs
phenicol, alkali, ruthenium(III) and periodate in turn increased with an increase in the concentration of
while keeping all other concentrations and conditions CHP (Table 1). It is observed that, as the CHP conꢀ
constant.
The oxidant DPC concentrations were varied in
the range of 5.0
10–6 to 5.0 10–5 mol/l, and fairly
constant kobs values indicated that order with respect to
DPC concentration was unity (Table 1). This was also
centration increases, the rate of reaction increases
with saturation: at [CHP]
≥ ×
2
10–3 mol/l, kobs almost
does not change, and the order with respect to CHP
×
×
concentration has changed from the first to zero order
(Fig. 3). In the lower concentration range, from 0.25
10–4 to 1.0 10–4 mol/l, the rate of oxidation shows
firstꢀorder dependence on [CHP]. At intermediate
concentrations of CHP, 1.0
10–4 to 5.0 10–4 mol/l,
the reaction order decreases, and, at the high concenꢀ
trations 2.0
10–3 mol/l, the rate of oxidation is
×
×
confirmed by linearity and almost parallel plots of log
vs. up to 80% completion of the reaction as shown in
Fig. 1.
A
t
×
×
The effect of chloramphenicol on the reaction rate
was studied at constant concentrations of alkali, DPC,
and periodate at the constant ionic strength of 0.10 mol/l.
The CHP concentration was varied in the range of
≥
×
independent of [CHP].
The effect of alkali concentration on the reaction
0.25
×
10–4 to 3.0
×
10–3 mol/l. The kobs values was studied in the range 0.01 to 0.1 mol/l at constant
KINETICS AND CATALYSIS Vol. 53
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68
HOSAHALLI et al.
−
Table 1. Effect of variation of [DPC], [CHP], [Ru(III)], [I ] and [OH–] on the oxidation of chloramphenicol by DPC
O4
at 25
°
C and
105, mol/l [CHP]
0.5
I = 0.10 mol/l
[DPC]
×
×
104, mol/l [OH–], mol/l
×
107, mol/l kobs 102, s–1
×
−
[I
O4
×
105, mol/l
[Ru(III)]
]
5.0
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.01
0.03
0.05
0.07
0.09
0.10
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
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
5.0
5.0
5.0
5.0
1.0
3.0
5.0
7.0
9.0
5.0
5.0
5.0
1.75
1.71
1.75
1.76
1.79
1.78
0.26
0.50
0.83
1.49
1.78
1.94
2.04
2.07
2.26
2.34
1.41
1.70
1.78
1.81
1.83
1.84
2.75
2.23
1.78
1.48
1.30
1.21
0.18
0.36
0.71
1.07
1.42
1.78
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
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
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
5.0
5.0
5.0
0.5
1.0
2.0
3.0
4.0
5.0
5.0
5.0
0.25
0.50
1.00
3.00
5.00
7.00
9.00
10.0
20.0
30.0
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
10.0
5.0
5.0
5.0
5.0
5.0
5.0
KINETICS AND CATALYSIS Vol. 53
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2012

KINETICS AND MECHANISM OF RUTHENIUM(III) CATALYZED OXIDATION
69
8 + log[Ru(III)]
0.8
0
0.4
1.2
1.6
2.0
2.0
2.0
1.6
1.2
0.8
0.4
0
1.6
1.2
0.8
0.4
0
0
1.0
2.0
[Ru(III)]
3.0
4.0
5.0
6.0
×
10–7, mol/l
5
Fig. 4. Effect of [Ru(III)] catalyst on the oxidation of CHP by DPC in aqueous alkaline medium at 25°C, [DPC] = 5.0 × 10 mol/l,
–4
–
[CHP] = 5.0 × 10 mol/l, [OH ] = 0.05 mol/l, and I = 0.10 mol/l.
concentrations of CPH, DPC, and periodate at conꢀ range 1.0
×
10–6 to 1.0
×
10–5 mol/l did not have any
stant ionic strength of 0.10 mol/l at 25 . The rate significant effect on reaction rate.
°
C
constants increased with an increase in the alkali conꢀ
centration (Table 1). The order was found to be less
than unity (0.31).
The participation of free radicals in the reaction
was examined as follows. The reaction mixture, to
which a known quantity of acrylonitrile monomer was
initially added, was kept for 2 h in an inert atmosphere.
On diluting the reaction mixture with methanol, a
white precipitate was formed indicating the participaꢀ
tion of free radicals. The blank experiments of either
DPC or Ru(III) alone with acrylonitrile did not
induce any polymerization under the same conditions
as those induced for reaction mixture. Initially added
acrylonitrile decreased the reaction rate also indicatꢀ
ing free radical intervention.
The effect of periodate was studied by varying the
periodate concentration from 1.0
×
10–5 to 1.0
×
10–4 mol/l keeping all other reactant concentrations
constant. It was found that the added periodate had a
retarding effect on the reaction rate (Table 1).
At constant oxidant, reductant, and alkali concenꢀ
tration of 5.0
tively, and = 0.10 mol/l, the catalyst ruthenium(III)
concentration was varied between 1.0
10–7 and
1.0
10–6 mol/l (Table 1). The order with respect to
×
10–5, 5.0
×
10–4 and 0.05 mol/l, respecꢀ
I
×
×
The kinetics was studied at four different temperaꢀ
tures under varying concentrations of CHP, alkali, and
periodate, keeping other conditions constant. The
ruthenium concentration was found to be unity. As
catalyst concentration increases the rate of the reacꢀ
tion also increases, and activity of catalyst reaches a
maximum after [Ru(III)]
≥
5.0
×
10–7 mol/l under the
above mentioned experimental conditions (Fig. 4).
Table 2. Thermodynamic activation parameters for the
Under condition used, the kobs for uncatalyzed reacꢀ Ru(III)ꢀcatalyzed oxidation of CHP by DPC in aqueous
tion is 1.44
10–3 s–1 and is negligible compared to the
×
alkaline medium with respect to the slow step of the scheme
mediated reaction (Table 1).
Δ
H# for
for Ki
kJ/mol
k
and
,
Δ
S# for
for Ki
J K–1 mol–1
k
and
,
Δ
G# for
for Ki
kJ/mol
k and
The effect of ionic strength was studied by varying
the potassium nitrate concentration from 0.01 to
1.0 mol/l at constant concentration of DPC, rutheꢀ
nium(III), CPH, periodate, and alkali. It was found
that increasing ionic strength had negligible effect on
the rate constant.
E
a for k,
Δ
H
Δ
S
Δ
G
,
kJ mol–1
14 0.5
12 0.2
26
–114
161
4
8
6
5
46 2.4
The effect of dielectric constant was studied by
–
–
–
2
–23 0.5
11.7 0.2
–23 0.4
varying the
tꢀbutanol–water (
v
/ ) content in the reacꢀ
v
tion mixture with all the other conditions being kept
constant. Decreasing the dielectric constant of the
medium had no effect on the rate of the reaction.
The externally added products,
hyde and copper(II) (CuSO4) in the concentration
–29 0.8
26 0.3
–135
161
pꢀnitrobenzaldeꢀ
KINETICS AND CATALYSIS Vol. 53
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70
HOSAHALLI et al.
reaction rate was found to increase with an increase in
temperature. The rate constants ( ) of the slow step of
Scheme were obtained from the slopes and intercepts of
[Ru(III)]/ obs vs. 1/[CHP],
obs vs. 1/[OH–], [Ru(III)]/
(a)
10–5, mol l–1
s
k
[Ru(III)]/kobs
8.0
×
1
2
k
k
and Ru(III)]/kobs vs. [H I 2−] plots (Figs. 5a–5c) at
O6
3
6.0
4.0
2.0
0
four different temperatures. The energy of activation
corresponding to these rate constants was evaluated
from the Arrhenius plot of logk vs. 1/T from which
3
4
otner activation parameters were calculated (Table 2).
It has been pointed out by Moelwyn–Hughes [24]
that, in presence of the catalyst, the uncatalyzed and
catalyzed reactions proceed simultaneously, so that
kobs
=
ku
+
K
C[catalyst]x .
(1)
Here kobs is the observed pseudoꢀfirst order rate conꢀ
stant in the presence of ruthenium(III) catalyst, ku is
the observed pseudoꢀfirst order rate constant for
uncatalyzed reaction, KC is the catalytic constant and
0
2.0
4.0
6.0
(b)
8.0
[1/CHP]
10.0
×
12.0
10–3, l/mol
[Ru(III)]/kobs × s
10–5, mol l–1
4.0
“
x
” is the order of the reaction with respect to
[Ru(III)]. In the present investigation, value is found
x
to be unity. Then the value of KC can be calculated
from Eq. (2)
1
4.0
4.0
4.0
4.0
2.0
0
KC = (kobs
–
ku)/[catalyst] =
kC/[catalyst],
(2)
2
3
4
where kC
=
kobs
–
ku. For the present experimental
condition and for the standard kinetic run, the KC was
found to be 3.26
×
104 mol–1 s–1 at 298 K.
l
The water soluble copper(III) periodate complex
[Cu(HIO6)2(OH)2]7– is reported [25]. However, in an
aqueous alkaline medium and at high pH range
employed in the study, periodate is unlikely to exist as
HI 4− that is evident from its involvement in the mulꢀ
O6
tiple equilibria (II)–(IV) [26–28], depending on the
pH of the solution.
0
2.0
4.0
6.0
8.0
10.0
12.0
2–
[H3IO6
]
×
10–3, mol/l
(c)
10–5, mol l–1
s
K
1
H5IO6
H4IO–6
H4IO–6 + H+
H3IO26– + H+
H2IO36– + H+
K1 = 5.1 ×
10–4, (II)
10–9, (III)
[Ru(III)]/kobs
6.0
×
K
2
K2 = 4.9 ×
1
2
4.5
3.0
1.5
0
K
H4IO26–
K
3 = 2.5 ×
10–12. (IV)
3
3
4
Periodic acid exists as H5IO6 in acid medium and as
−
H I
near pH 7. Thus, under the conditions
employed in alkaline medium, main species are
O6
4
2−
3−
expected to be H I
O6
and H I Thus, at pH
O6 .
0
2.0
4.0
6.0
8.0
1/[OH]
10.0
10–1, l/mol
12.0
3
2
employed in this study, the soluble copper(III) perioꢀ
date complex exist as diperiodatocuprate(III),
[Cu(OH)2(H3IO6)2]2–, a conclusion also supported by
the literature [27, 28].
The reaction between the diperiodatocuprate(III)
complex and CHP in alkaline medium has stiochiomꢀ
etry 1 : 2 (CHP : DPC) with first order dependence on
×
Fig. 5. Plots of [Ru(III)]/
k
vs. 1/[CHP] (a), [H IO62−
]
obs
3
–
(b), and 1/[OH ] (c) at different temperatures, K: 288 (
1),
298 (2), 308 (3), and 318 (4) (conditions as in Table 1).
KINETICS AND CATALYSIS Vol. 53
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KINETICS AND MECHANISM OF RUTHENIUM(III) CATALYZED OXIDATION
71
the DPC and ruthemium(III) concentrations and less
than the first order in CHP and alkali concentrations,
a negative fractional order dependence on the perioꢀ
date concentration. No effect of added products was
observed. Based on the experimental results, a mechꢀ
anism is proposed for which all the observed orders in
respect to each constituent such as [oxidant], [reducꢀ
[Cu(H3IO6)2]– + [OH–]
[Cu(H2IO6)(H3IO6)]2– + H2O,
K
4
[Cu(H2IO6)(H3IO6)]2– + 2H2O
[Cu(H2IO6)(H2O)2] + [H3IO6]2–,
K
5
H
C
NHCOCHCl2
C CH2OH
tant], [RuIII], [OH–] and [
I
−] may be well accommoꢀ
O4
O2N
dated. Lister [29] proposed three forms of copper(III)
periodate in alkaline medium as diperiodatocuꢀ
prate(III) (DPC), monoperiodatocuprate(III) (MPC)
and tetraperiodatocuprate(III). The latter is ruled out,
OH H
K
6
+ [Ru(H2O)5OH]2+
Complex C1 + H2O,
Complex C1 + [Cu(H2IO6)(H2O)2]
as its equilibrium constant is 8.0 × °C.
10–11 mol/l at 40
In the present kinetic study, DPC and MPC are to be
considered as active forms of copper(III) periodate
complex. It may be expected that lower periodate
complex such as MPC is more important in the reacꢀ
tion than DPC. The increase in the rate with an
increase in alkali concentration suggests that equilibia
of copper(III) periodate complexes are possible as in
Eqs. (V) and (VI).
H
C
NHCOCHCl2
k
slow
O2N
C CH2OH
H
+
O
•
+ [Ru(H2O)5OH]2+ + H+ + Cu2+ + H2IO63– + 2H2O,
NO2
H
C
NHCOCHCl2
C CH2OH
H
fast
O2
N
O
•
[Cu(H3IO6) ] + [OH–]
–
O
C H
2
NHCOCHCl2
(V)
K
[Cu(H2IO6)(H3IO6)]2– + H2O,
4
•
+
,
C CH2OH
H
NHCOCHCl2
[Cu(H2IO6)(H3IO6)]2– + 2H2O
[Cu(H2IO6)(H2O)2] + [H3IO6]
+ [Cu(H2IO6)(H2O)2] + OH–
•
C CH2OH
H
(VI)
K
2–
5
.
NHCOCHCl2
fast
+ Cu2+ + H2IO63– + 2H2O,
HO C CH2OH
H
The reaction order less than unit in respect to
[CHP] presumably results from the formation of a
complex C1 between the ruthenium(III) species and
CHP prior to the formation of the products. The
ruthenium(III) catalyst forms the complex with the
substrate, which enhances the reducing property of
the substrate. This complex C1 reacts with one moleꢀ
cule of MPC in the fast step to give a free radical interꢀ
mediate species from CHP, Cu(III) being reduced to
Cu(II) by accepting an electron from O–H bond
cleavage and ruthenium(III) being regenerated. The
C–C bond in the free radical cleaves in a subsequent
NHCOCHCl2
NH2
fast
+ H2O
HO C CH2OH
H
HO C CH2OH
H
+ Cl2CHCOOH,
fast
H+ + OH–
H2O.
Scheme.
The probable structure of the complex C1 is given by:
2+
H
C
O
NHCOCHCl2
fast step to give pꢀnitrobenzaldehyde and another free
radical. In the next fast step, the free radical thus
formed reacts with another molecule of MPC in the
presence of OH– to form 2,2ꢀdichloroꢀNꢀ(1,2ꢀdihyꢀ
droxyethyl)acetamide. In the further fast steps 2,2ꢀ
dichloroꢀNꢀ(1,2ꢀdihydroxyethyl)acetamide undergoes
hydrolysis to give another product 2ꢀamino~2ꢀhydroxꢀ
yethanol and dichloroacetic acid. The detailed mechꢀ
anism for the oxidation of CHP by diperiodatocuꢀ
prate(III) is presented in scheme.
O2N
C CH2OH
H
H2O
H2O
OH
Ru
OH2
OH2
.
Complex C1 shows that hydroxyl group oxygen was
coordinated to the metal complex. Since Scheme is in
accordance with the generally well accepted principle
KINETICS AND CATALYSIS Vol. 53
No. 1
2012

72
HOSAHALLI et al.
2–
of nonꢀcomplementary oxidations taking place in
sequence of oneꢀelectron steps, the reaction between
substrate and oxidant would afford a radical intermeꢀ
diate. A free radical scavenging experiment revealed
the possibility. This type of radical intermediate has
also been observed in earlier work [30]. A bathochroꢀ
mic shift of about 6 nm from 271 to 277 nm in the
spectra of CHP observed confirms the complex forꢀ
mation between ruthenium(III) and CHP in their
mixture. The Michaelis–Menten plot also proves the
complex formation between the catalyst and substrate
and explains the complex dependence on [CHP].
Such type of complex formation between a substrate
and a catalyst has been observed in literature [31, 32].
The rate law (11) for the intermediate concentrations
of DPC can be derived from scheme as follows.
According to scheme,
[DPC]t = [DPC]f + [Cu(H2IO6)(H2IO6)]
[Cu(H2IO6)(H2O)2] = [DPC]f
+
(4)
⎫
2–
2–
–
–
⎧
[H3IO6 ] + K4[H3IO6 ][OH ] + K4K5[OH ]
× ꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀ .
⎨
⎬
[H3IO26–]
⎩
⎭
Therefore [DPC]f is given by
[DPC]f
[DPC]t[H3IO26–]
ꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀ .
(5)
=
[H3IO26–] + K4[H3IO26–][OH–] + K4K5[OH–]
Since concentrations of ruthenium(III) used are sevꢀ
eral orders of magnitude lower than [CHP],
d[DPC]
[CHP]t = [CHP]f.
(6)
Rate = –ꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀ = k[Cu(H2IO6)(H2O)2[C1]
dt
The free [Ru(IIl)] can be written as
2–
kK5K6[Cu(H2IO6)(H3IO6)] [CHP][Ru(III)]
Ru III
(
)
]
[
=
ꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀ
f
[H3IO26–]
[Ru(III)] =
f
(7)
Ru III
=
.
(
)
]
[
(3)
f
1 + K6 CHP
[
]
kK K K [CHP][Ru(III)][DPC] [OH–]
In the same manner, due to the low concentrations
f
4
5
6
f
=
ꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀ ,
[H3IO26–]
of [DPC] and [H I 2−] in comparison to [OH–
]
O6
3
[OH–]f = [OH–]t.
(8)
where [DPC]t and [DPC]f refer to total and free DPC
concentrations. The total concentration of DPC, Substituting Eqs. (5)–(8) in Eq. (3) and omitting t and
[DPC]t,is given by
f, we get
kK4K5K6[CHP][DPC][OH–][Ru(III)]
Rate = ꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀ, or
[H3IO26–] + K4[H3IO62–][OH–] + K4K5[OH–] + K4K5K6[OH–][CHP]
(9)
kK4K5K6[CHP][DPC][OH–][Ru(III)]
[H3IO26–] + K4[H3IO62–][OH–] + K4K5[OH–] + K4K5K6[OH–][CHP]
Rate
[DPC]
ꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀ = kobs = ꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀ.
(10)
The rate law (10) can be transformed to the following are in good agreement with those given in the literaꢀ
form, which is suitable for verification.
ture [33]. The equilibrium constant K4 is far greater
than K5. This may be attributed to the greater tendency
of DPC to undergo hydrolysis compared to the dissoꢀ
ciation of hydrolyzed species in alkaline medium.
The thermodynamic quantities for the different
equilibrium steps of the scheme can be evaluated as
follows. The [CHP] and [OH–] (as in Table 1) were
varied at four different temperatures. The plots of
[H3IO26–]
[Ru(III)]
ꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀ = ꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀ
kK4K5K6[OH–][CHP]
kobs
(11)
[H3IO26–]
1
1
k
+
ꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀ + ꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀꢀ + ꢀ.
kK5K6[CHP] kK6[CHP]
[Ru(III)]/k kobs vs. 1/[CHP],
obs vs. 1/[OH–], [Ru(III)]/
According to Eq. (11), other conditions being conꢀ
and [Ru(III)]/kobs vs. [H I 2−] are linear. From the
stant, plots of[Ru(III)]/
obs vs. 1/[OH–], [Ru(III)]/kobs
k
O6
3
slopes and intercepts, the values of K4 were calculated
at different temperatures, and these values are given in
vs. 1/[CHP], and Ru(III)]/kobs vs. [H I 2−] should be
O6
3
linear and found to be so (Fig. 5). From the slopes and
Table 3. vant Hoff dependences (logK4 vs. 1/
plotted, and the values of the enthalpy of reaction
the entropy of reaction , and the free energy of reacꢀ
tion were calculated (Table 2). A comparison of the
these values with those obtained from the slow step of
T) were
intercepts of these plots the following values of K4
K6 and at 25 were calculated: 2.11
10–1 mol–1,
10–2 mol/l, 6.43 103 l mol–1 s–1, and
104 s–1, respectively. The values of K4
K5 and K6
, K5,
Δ
H,
k
°
C
×
1
Δ
S
1.62
4.98
×
×
Δ
G
×
,
KINETICS AND CATALYSIS Vol. 53
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2012

KINETICS AND MECHANISM OF RUTHENIUM(III) CATALYZED OXIDATION
73
Table 3. The values of K4, K5 and K6 calculated at different temperatures
Temperature, K
k
×
10–4, s–1
K
4, l/mol
K5
×
102, mol/l
K6 ×
10–3, l/mol
288
298
308
318
3.88
4.98
5.83
6.97
0.06 0.01
0.21 0.01
0.37 0.01
0.41 0.02
2.12 0.01
1.62 0.01
1.12 0.02
0.65 0.01
4.1 0.2
6.3 0.2
7.7 0.1
12.3 0.1
2. Rao, P.J.P., Sethuram, B., and Navaneeth Rao, T.,
J. Indian Chem. Soc., 1990, vol. 67, p. 101.
3. Murthy, C.P., Sethuram, B., and Navanaeeth Rao, T.,
Z. Phys. Chem., 1986, vol. 267, p. 1212.
4. Niu, W., Zhu, Y., Hu, K., Tong, C., and Yang, H., Int.
J. Chem. Kinet., 1996, vol. 28, p. 899.
5. Rozovoskii, G.I., Misyavichyus, A.K., and Prokoꢀ
pchik, A.Y., Kinet. Catal., 1975, vol. 16, p. 337.
6. Ram, ReddyM.G., Sethuram, B., and Navaneeth
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 fast and
occurs with the low activation energy [34]. In the same
manner, K5 and K6 values were calculated at different
temperatures (Table 3), and the corresponding values
of the thermodynamic parameters are given in Table 2.
The moderate values of
mediate complex is more ordered than the reactants
[35]. The value of
S# within the range of radical reacꢀ
Δ
S# suggest that the interꢀ
Rao, T., Indian J. Chem., Sect. A, 1978, vol. 16, p. 331.
Δ
tion has been ascribed to the nature of electron pairing
and unpairing processes and to the loss of degrees of
freedom formerly available to the reactants upon the
formation of rigid transition state [36]. The observed
modest enthalpy of activation and relatively low value
of the entropy of activation as well as a higher rate conꢀ
stant of slow step indicate that the oxidation presumꢀ
ably occurs via innerꢀsphere mechanism. The concluꢀ
sion is supported by earlier observation [37]. The negꢀ
ligible effect of ionic strength explains the involvement
of neutral species in scheme. However, it is difficult to
interpret the effect of dielectric constant in view of
various ions involved. The catalyst ruthenium(III)
forms the complex with substrate, which enhances the
reducing property of substrate.
7. Karlin, K.D. and Gulineh, Y., Progress in Inorganic
Chemistry, New York: Wiley, 1997, vol. 35, p. 220.
8. Tolman, W.B., Acc. Chem. Res., 1997, vol. 30, p. 227.
9. Kovat, Z., Acta Chim. Hung., 1960, vol. 22, p. 313.
10. Kitajima, K.N. and Moroꢀoka, Y., Chem. Rev., 1994,
vol. 94, p. 737.
11. Karlin, K., Kaderli, S., and Zuberbuhler, A.D., Acc.
Chem. Res., 1997, vol. 30, p. 139.
12. Piere, J.L., Chem. Soc. Rev., 2000, vol. 29, p. 251.
13. Solomon, E.I., Chen, P., Metz, M., Lee, S.K., and
Palmer, A.E., Angew. Chem., Int. Ed. Engl., 2001,
vol. 40, p. 4570.
14. Halcrow, M.A., Angew. Chem., Int. Ed. Engl., 2001,
vol. 4, p. 816.
15. Peisach, J., Alsen, P., and Blumberg, W.E., The Bioꢀ
chemistry of Copper, New York: Academic, 1966, p. 49.
16. Meenakshisundaram, S. and Sathiyendiran, V., J. Chem.
Res., 2000, vol. 10, p. 458.
Among various species of DPC in alkaline
medium, monodiperiodatocuprate(III) (MPC)
(
Cu(H2IO6)(H2O)2) is considered as active species for
the reaction under study. The active species of rutheꢀ
nium(III) is found to be [RuH2O)5OH]2+. The cataꢀ
lytic constants and the activation parameters with refꢀ
erence to catalyst were computed. The results demonꢀ
strate that the role of pH in the reaction medium is
crucial. Rate constant of the slow step and other equiꢀ
librium constants involved in the mechanism evaluꢀ
ated, and activation parameters with respect to slow
step of the reaction were computed. The overall mechꢀ
anistic sequence described here is consistent with
product, mechanistic, and kinetic studies.
17. Kambo, N. and Upadhyaya, S.K., Transition Met.
Chem., 2000, vol. 25, p. 461.
18. Reddy, C.S. and Vijaykumar, T., Indian J. Chem., 1995,
vol. 34A, p. 615.
19. Murthy, C.P., Sethuram, B., and Navaneeth Rao, T.,
Z. Phys. Chem., 1981, vol. 262, p. 336.
20. Jeffery, G.H., Bassett, J., Mendham, J., and
Denny, R.C., Vogel’s Textbook of Quantitative Chemical
Analysis, New York: Longman, 1996, p. 455.
21. Panigrahi, G.P. and Misro, P.K., Indian J. Chem., 1978,
vol. 16A, p. 201.
22. Randrerath, K., Thin Layer Chromartography, New
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KINETICS AND CATALYSIS Vol. 53
No. 1
2012

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KINETICS AND CATALYSIS Vol. 53
No. 1
2012