Mechanism of oxidation of ketorolac by hexacyanoferrate(III) in aqueous alkali: A thermodynamics and kinetics study

Article abstract of DOI:10.1007/s10953-014-0177-0

The kinetics of the oxidation of ketorolac by hexacyanoferrate(III) (HCF) in aqueous alkaline medium at a constant ionic strength of 0.75 mol·dm^-3 was studied spectrophotometrically at 300 K. A plausible mechanism was proposed and the rate law was derived. The mechanism of oxidation of ketorolac (KET) in alkaline medium has been shown to proceed via a KET-HCF complex, which decomposes in a slow step followed by other fast steps to give the products. The main oxidative product was identified as (2,3-dihydro-1-hydroxy-1H-pyrrolizin-5-yl-)(phenyl)methanone and is characterized by its LC-ESI-MS spectrum. Thermodynamic parameters of various equilibria of the mechanism were calculated and activation parameters ΔH ^A, ΔS ^A, ΔG ^A and log₁₀ A were found to be 29.9 kJ·mol^-1, -220 J·K ^-1·mol^-1, 96 kJ·mol^-1 and 2.70 respectively.

Full text of DOI:10.1007/s10953-014-0177-0

J Solution Chem (2014) 43:916–929
DOI 10.1007/s10953-014-0177-0
Mechanism of Oxidation of Ketorolac
by Hexacyanoferrate(III) in Aqueous Alkali:
A Thermodynamics and Kinetics Study
Seema S. Badi Suresh M. Tuwar
•
Received: 22 December 2013 / Accepted: 7 February 2014 / Published online: 21 May 2014
Ó Springer Science+Business Media New York 2014
Abstract The kinetics of the oxidation of ketorolac by hexacyanoferrate(III) (HCF) in
-3
aqueous alkaline medium at a constant ionic strength of 0.75 molꢀdm was studied
spectrophotometrically at 300 K. A plausible mechanism was proposed and the rate law
was derived. The mechanism of oxidation of ketorolac (KET) in alkaline medium has been
shown to proceed via a KET-HCF complex, which decomposes in a slow step followed by
other fast steps to give the products. The main oxidative product was identified as (2,
3-dihydro-1-hydroxy-1H-pyrrolizin-5-yl-)(phenyl)methanone and is characterized by its
LC–ESI–MS spectrum. Thermodynamic parameters of various equilibria of the mecha-
=
=
=
nism were calculated and activation parameters DH , DS , DG and log10 A were found
-1
-1
-1
-1
to be 29.9 kJꢀmol , -220 JꢀK ꢀmol , 96 kJꢀmol and 2.70 respectively.
Keywords Hexacyanoferrate(III) ꢀ Ketorolac ꢀ Oxidation ꢀ Kinetics ꢀ
Thermodynamics ꢀ Mechanism
1
Introduction
Ketorolac (KET, systematic name: 5-benzoyl-2,3-dihydro-1H-pyrrolizine-1-carboxylic
acid) is a non-steroidal, anti-inflammatory drug (NSAID) with potent analgesic and anti-
pyretic activity that exerts its actions through the inhibition of prostaglandin synthesis by
acting as a non-selective competitive blocker of the cyclooxygenase pathway of arachid-
onate metabolism [1, 2]. Prostaglandin is involved in multiple homeostatic processes; KET
like other NSAIDS may induce adverse events involving the gastrointestinal, renal and
blood clotting systems. It may exert its analgesic activity via its effects on the central
nervous system. Both racemic forms of this compound are known, where the S antipode of
S. S. Badi ꢀ S. M. Tuwar (&)
Department of Chemistry, Karnatak Science College, Dharwad 580 001, Karnataka, India
e-mail: sureshtuwar@yahoo.co.in; sm.tuwar@gmail.com
1
23

J Solution Chem (2014) 43:916–929
917
KET is more potent and causes fewer side effects [3]. KET is also used to treat eye pain
and to relieve the itchiness and burning of seasonal allergies [4, 5].
Many transition and non-transition metal ions, in their complexed forms, act as good
oxidants in acidic, basic or neutral media. However, their oxidation capacity depends on
the redox potential, which is varied by changing the pH. Hexacyanoferrate(III) (HCF), one
such oxidant in alkaline medium is well understood [6–12], particularly its oxidative
capacity in oxidation of inorganic and organic compounds. The reduction potential of the
3
-
4-
[
Fe(CN)6] /[Fe(CN)6] couple remains essentially constant at around ?0.41 V in the pH
3
-
range 4–13, with oxidized, reduced or protonated forms [13]. Therefore, [Fe(CN)6] is
considered as a mild oxidant. Only at higher acidities, due to the association with protons,
-
the species H[Fe(CN)6] , H [Fe(CN) ] and even H [Fe(CN) ] can be formed [14] and
2
-
2
6
3
6
the potential increases from ?0.56 to ?0.93 V in the concentration range
-
3
?
0
.1–8.0 molꢀdm [H ]. This indicates that HCF is a moderate oxidant in aqueous alkali
and strong at higher acidities. It is a single equivalent and stable oxidant, and oxidizes the
organic compounds through an outer-sphere mechanism [15, 16]. Oxidation of organic
compounds by such a stable oxidant adds less error to the experimental results and data can
be analyzed meticulously to establish the reaction path.
It is also well known that HCF is an hydrogen abstracting agent [6] for most easily
oxidizable substrates. Hence, it is used as an interceptor for free radicals and as an electron
abstracting reagent [8, 10] both in alkali and neutral media [11]. HCF is employed to
oxidize a verity of organic substrates including thiols, thiol acids and thioamides. The key
step in the oxidation is an electron transfer, generally acknowledged to proceed by the
formation of substrate-oxidant ion pair or outer sphere precursor complex in a pre equi-
librium step. Formation of such species is expected to bring about changes into the
vibrational movements [17].
Since the KET is non-sulfural compound, the present study is to understand whether its
redox chemistry with HCF proceeds with the formation of an substrate–oxidant ion pair or
outer sphere precursor complex, in a pre equilibrium step. Hence, the title reaction is
undertaken to study the mechanism of oxidation of KET.
2
Experimental
2
.1 Chemicals
All chemicals used were reagent grade and double distilled water was used throughout the
study. The KET was obtained from Raptakos Brett and Co., Microlabs Ltd. KLAB,
Mumbai, India as free sample with 98.8 % purity. Further, its purity was checked by its
m.p. and TLC for single spot. A stock solution of KET was prepared by dissolving the
required quantity of sample. The kinetics runs were checked with aqueous solutions of
KET with different ageing of the solution to understand the stability of KET solutions. It
was found that there was no difference between fresh and aged solutions. However, fresh
solutions were used in all kinetic runs. Aqueous solutions of KOH and KCl were used to
-
maintain required [OH ] and ionic strength, respectively. A solution of HCF was prepared
by dissolving K [Fe(CN) ] (BDH) in water and was standardized iodometrically in the
3
6
4
presence of ZnSO to avoid the reverse reaction of [Fe(CN) ] [18]. K [Fe(CN) ] was
-
4
6
4
6
4
-
used to prepare [Fe(CN)6] to study the product effect. Earlier it was estimated quanti-
3
tatively by oxidizing it into [Fe(CN)6]
-
.
123

918
J Solution Chem (2014) 43:916–929
2
.2 Kinetic Measurements
The oxidation of KET by HCF was carried out under pseudo-first order conditions, where
KET] [ [HCF] at 300 ± 1 K unless otherwise specified. The reaction was initiated by
[
mixing the required quantities of previously thermostated solutions of KET and HCF,
which also contained the necessary quantities of KOH and KCl to maintain the required
alkalinity and ionic strength. The course of the reaction was followed by measuring the
decrease in absorbance of HCF in the reaction mixture in a 1 cm quartz cell in the
thermostated compartment of a Specord-200 plus spectrophotometer, set up with a Peltier
accessory to control temperature. The absorption measurements were performed at its
absorbance maximum, 420 nm, as a function of time. Earlier, it was verified that there is
negligible interference from other species present in the reaction mixture at this wave-
length. The obedience of absorbance by HCF to Beer’s law at the wavelength 420 nm was
verified earlier and the molar absorbance coefficient ‘e’ was found to be 1,050 (± 20)
3
-1
-1
dm ꢀmol ꢀcm . First order rate constants, kobs, were calculated from plots, log [HCF]
10
against time. Excellent first order plots were obtained up to 90 % completion. The results
are reproducible within ± 3 %.
2
.3 Stoichiometry and Product Analysis
-
Eight different sets of concentrations of KET and HCF were kept at constant [OH ] and
ionic strength in an inert atmosphere at 300 K for over 24 h in a closed vessel, under an
inert atmosphere, for completion of the reaction. Unreacted HCF was estimated spectro-
photometrically by measuring its absorbance at 420 nm. Further, the experiments were
repeated at higher concentrations for titrimetric analysis. Under such conditions, and when
HCF is in excess of KET, the unreacted HCF was estimated iodometrically.
The results indicated that two moles of HCF were consumed by one mole of KET as in
Eq. 1.
O
C
O
C
N
N
COOH
OH
3
-
-
+
2[Fe(CN) ] + 2OH
6
4-
+
2[Fe(CN) ] + H O + CO
6
2
2
1
The oxidative product of KET was identified as (2,3-dihydro-1-hydroxy-1H-pyrrolizin-
-yl-)(phenyl)methanone and was ascertained as follows. After completion of the reaction,
5
the reaction mixture was subjected to TLC for separation of the constituents with reference
to the KET. Iodine spray showed a single spot, indicating the presence of only one product.
Further, the solution was subjected for LC–ESI–MS. The reaction mixture, after com-
pleting the reaction, was treated with 50 % methanol followed by acidification with HCl
and 3 % acetonitrile and 1 % formic acid to make the solution suitable for a positive ion
mode for the LC–ESI–MS analysis. The solution was subjected to LC–ESI–MS analysis at
-
1
the rate of 5 lLꢀmin , with retention time 0.51–0.98 s at the applied voltage of 30 kV,
using a glass micro syringe. Nitrogen gas was used as nebulizer. The LC–ESI–MS analysis
1
23

J Solution Chem (2014) 43:916–929
919
of reaction at positive ionic mode indicated the presence of a product with a molecular ion
-
1
-1
peak, mꢀz at 225 (mꢀz 224 ± 3) as was expected for (2,3-dihydro-1-hydroxy-1H-
pyrrolizin-5-yl-)(phenyl)methanone (Fig. 1). Further, the retention of ketone and formation
of new OH due to decarboxylation were confirmed by a spot test [19].
3
Results and Discussion
3
.1 Reaction Order
The orders, with respect to [KET] and [alkali], were found from the graph of log10 kobs
against log10 (concentration) plots by varying the concentration of reductant and alkali in
turn while keeping the others constant.
3
.2 Effect of [HCF]
The effect of [HCF] on the reaction rate was studied by varying the [HCF] from
-
5
-4
-3
8
strength at a constant temperature (Table 1). It was observed that slope of the plot of log
.0 9 10 to 10.0 9 10 molꢀdm at constant concentrations of KET, alkali and ionic
1
0
(conc.) against time remained constant for most of the [HCF] values. This was also
confirmed from the linearity of plots of log10 (absorbance) against time to about 80 %
completion of the reaction. Hence, the order in [HCF] was considered as being unity.
3
.3 Effect of [KET]
The effect of [KET] on rate of reaction was studied at different temperatures in the
-
3
-2
-3
concentration range 1.0 9 10
to 1.0 9 10 molꢀdm
at all other reaction
8
6
4
2
0
0
0
0
0
O
OH
N
225
Mol. mass = 225
0
50
100
150
m/z
200
250
Fig. 1 LC-ESI–MS spectrum of the product, (2,3-dihydro-1-hydroxy-1H-pyrrolizin-5-yl-)(phenyl)metha-
none resulted due to the oxidation of ketorolac by hexacyanoferrate(III) in aqueous alkaline medium
123

920
J Solution Chem (2014) 43:916–929
-
Table 1 Effect of variations of [HCF], [KET] and [OH ] on oxidation of ketorolac by hexacyanofer-
-
3
rate(III) in aqueous alkaline medium at 300 K, I = 0.75 molꢀdm
4
0 9 [HCF]
3
10 9 [KET]
-
[OH ]
3
-1
1
10 9 kobs (s
)
-
3
)
-3
-3
)
(molꢀdm
(molꢀdm
)
(molꢀdm
a
b
Exptl.
Calc.
0
1
2
4
6
8
1
4
4
4
4
4
4
4
4
4
4
4
4
4
4
4
4
4
a
.8
.0
.0
.0
.0
.0
0
4.0
4.0
4.0
4.0
4.0
4.0
4.0
1.0
2.0
4.0
6.0
8.0
10
0.50
0.50
0.50
0.50
0.50
0.50
0.50
0.50
0.50
0.50
0.50
0.50
0.50
0.25
0.30
0.35
0.40
0.45
0.50
0.55
0.60
0.65
0.70
0.75
3.00
2.90
3.09
3.21
3.20
3.17
3.13
1.07
2.53
3.21
4.70
6.16
7.67
1.76
2.11
2.31
2.58
2.85
3.21
3.38
3.69
3.93
4.20
4.36
–
–
–
–
–
–
–
.0
.0
.0
.0
.0
.0
.0
.0
.0
.0
.0
.0
.0
.0
.0
.0
.0
1.65
3.07
5.40
7.23
8.71
9.93
3.25
3.75
4.21
4.63
5.03
5.40
5.75
6.07
6.37
6.66
6.93
4.0
4.0
4.0
4.0
4.0
4.0
4.0
4.0
4.0
4.0
4.0
Experimental
b
-2
-1
3
-1
Calculated:
k
obs were calculated by using k = 6.66 9 10
-1
s
,
K
1
= 1.02 dm ꢀmol
and
3
K
2
= 78.7 dm ꢀmol at 300 K in Eq. 7
concentration constant (Table 2).The kobs values increased with increase in [KET]. The
order in [KET] from the plot of log10 kobs against log10 [KET] was found to be a positive
fractional order but decreased marginally from 0.70 to 0.45 in the temperature range
3
00–320 K.
-
.4 Effect of [OH ]
3
-
The influence of [OH ] on the rate of reaction was studied by varying [KOH] from 0.25 to
-
3
0
range 300 to 320 K. It was observed that rates increased with increase in [OH ] (Table 3).
.75 molꢀdm while keeping all other reaction concentration constant in the temperature
-
-
The first order plots were linear for all the [OH ] in different temperatures, and the order
-
from the plot of log10 kobs against log [OH ] was found to be positive fractional, but the
10
order decreases from 0.82 to 0.56 as temperature increases, in the range 300–320 K, at a
-
3
fixed ionic strength of 0.75 molꢀdm
.
1
23

J Solution Chem (2014) 43:916–929
921
Table 2 Effect of variation of [KET] on oxidation of ketorolac by hexacyanoferrate(III) in aqueous
-4
-3
-
-3
alkaline medium at different temperatures: [HCF] = 4.0 9 10 molꢀdm , [OH ] = 0.50 molꢀdm , and I =
-
3
0
.75 molꢀdm
3
0 9 [KET]
3
-1
1
10 9 kobs (s
)
-
3
)
(molꢀdm
a
b
Exptl.
Calc.
for temp
00 K
3
00 K
305 K
310 K
315 K
320 K
3
1
2
4
6
8
1
.0
.0
.0
.0
.0
0
1.07
2.53
3.21
4.70
6.16
7.67
0.70
2.10
3.18
5.32
6.95
8.50
9.47
0.50
2.19
3.67
6.32
8.27
9.54
10.6
0.48
2.85
4.53
6.23
9.10
10.6
13.4
0.47
3.06
5.14
7.46
10.7
12.2
13.5
0.45
1.65
3.07
5.40
7.23
8.71
9.93
0.70
order
a
Experimental
b
-2
-1
3
-1
Calculated:
k
obs were calculated by using k = 6.66 9 10
-1
s
,
K
1
= 1.02 dm ꢀmol
and
3
K
2
= 78.7 dm ꢀmol at 300 K in Eq. 7
-
Table 3 Effect of variation of [OH ] on oxidation of ketorolac by hexacyanoferrate(III) in aqueous
-
4
-3
-3
-3
alkaline medium at different temperatures: [HCF] = 4.0 9 10 molꢀdm , [KET] = 4.0 9 10 molꢀdm
,
-3
and I = 0.75 molꢀdm
-
OH ]
3
-1
[
10 9 kobs (s
)
-
3
)
(molꢀdm
a
b
Exptl.
Calc. for
temp 300 K
300 K
305 K
310 K
315 K
320 K
0
0
0
0
0
0
0
0
0
0
0
.25
.30
.35
.40
.45
.50
.55
.60
.65
.70
.75
1.76
2.11
2.31
2.58
2.85
3.21
3.38
3.69
3.93
4.20
4.36
0.82
2.70
2.92
3.26
3.50
3.81
4.09
4.24
4.67
3.97
5.08
5.40
0.65
2.92
3.38
3.85
4.22
4.56
5.04
5.16
5.30
5.52
5.42
5.97
0.62
3.54
4.10
4.01
4.53
4.47
4.97
5.34
4.95
5.51
5.96
6.08
0.58
4.33
4.78
4.68
4.99
6.58
6.98
6.91
7.12
7.30
7.65
8.35
0.56
3.25
3.75
4.21
4.63
5.03
5.40
5.75
6.07
6.37
6.66
6.93
0.82
order
a
Experimental
b
-2 -1
3
-1
Calculated:
k
obs were calculated by using k = 6.66 9 10 ꢀs
,
K
1
= 1.02 dm ꢀmol
and
3
-1
K
2
= 78.7 dm ꢀmol at 300 K in Eq. 7
3
.5 Effect of Initially Added Product
4-
The effect of initially added [Fe(CN) ] on the rate of reaction was studied by varying
6
4
-
-5
-4
K [Fe(CN) ]
4
solution concentrations in the range of 8.0 9 10
3-
to 8.0 9 10
6
-3
-
molꢀdm at 300 K at constant [Fe(CN)6] , [KET], [OH ] and ionic strength. The added
product did not alter the rate of reaction.
123

922
J Solution Chem (2014) 43:916–929
3
.6 Effect of Ionic Strength and Dielectric Constant of the Medium
-
3
The ionic strength of the reaction medium was varied between 0.5 and 1.0 molꢀdm with
potassium chloride at constant concentrations of alkali, oxidant and reductant at 300 K. It
was found that as ionic strength increases, the rate of reaction increases. A plot of log10 kobs
against HI was linear with positive slope (0.7).
The relative permittivity (D) effect was studied by varying the t-butyl alcohol–water
content in the reaction mixture with all other conditions being kept constant. The D values
were computed [20] from the values of the pure liquids. In the reaction, as D increases the
-1
kobs values increased. The graph of log10 kobs against D was found to be linear with a
negative slope. Earlier the reaction between t-butyl alcohol and oxidant was studied. It was
observed that there was no reaction between solvent and oxidant.
3
.7 Effect of Temperature
The effect of temperature on the reaction rate was studied at various temperatures at
constant concentrations of oxidant and reductant, with other reaction conditions being kept
constant. The kobs values were found increase with temperature (Table 4). The activation
energy was calculated from the slope of Arrhenius plot and, by using this, the other
activation parameters were calculated and tabulated in Table 5.
Since the reaction follows a multistep mechanism, the reaction constants: equilibrium
constants and rate constant of the slow step are involved. In order to arrive at the values of
these constants, the reaction was studied at 300, 305, 310, 315 and 320 K by varying
-
[
KET] and [OH ] in turn while other concentrations were kept constant (Tables 2 and 3).
The rate constants (k) of the slow step of Scheme 1 were obtained from the intercept of
-
1
- -1
against [OH ] for the temperatures 300, 305, 310, 315 and 320 K. The values
(
‘
kobs
)
k’ calculated for slow step decrease with increasing temperature. Hence, activation
parameters for the slow step are not calculated. This may be due to the instability of the
adduct at higher temperatures (discussed elsewhere).
The equilibrium constants K and K of the scheme were determined from the slopes
1
2
-
1
-1
- -1
and intercepts of the plots of (kobs
)
against [KET] and [OH ] for different tem-
peratures (Table 6). From the van’t Hoff plot, the thermodynamic parameters DH*, DS*
and DG* (under the reaction conditions and temperature) for the formation of KET anion,
and complex between HCF and KET anion, were determined (Table 7).
3
.8 Polymerization Study
As HCF is a single equivalent oxidant, the oxidation of organic compounds by such an
oxidant could lead to free radical generation. Hence, to test for the possible intervention of
free radicals, the reaction mixture was mixed with acrylonitrile monomer and kept for
2
4 h. On dilution with methanol, a white precipitate of polymer formed, indicating that the
reaction occurs with the intervention of free radicals. The experiment of either HCF or
KET with acrylonitrile alone did not induce the polymerization under similar conditions as
those induced with the reaction mixture. This was also evidenced by decrease in the kobs
values in the presence of initially added acrylonitrile.
When the HCF reaction order in the oxidation of organic substrates is unity or close to
unity, both in oxidant and reductant species, the rate determining step is the transfer of one
4
electron from substrate to oxidant to produce [Fe(CN )] . When the product inhibits the
-
6
overall reaction, this step becomes reversible [11, 21]. Both possibilities are unlikely to be
1
23

J Solution Chem (2014) 43:916–929
923
Table 4 Effect of temperature on the oxidation of ketorolac by hexacyanoferrate(III) in aqueous alkaline
-
4
-3
-3
-3
-
-3
medium: [HCF] = 4.0 9 10 molꢀdm , [KET] = 4.0 9 10 molꢀdm , [OH ] = 0.50 molꢀdm , and I =
-3
0
.75 molꢀdm
3 -1
10 9 kobs (s ) (overall reaction)
2
-1
)
Temperature (K)
10 9 kobs (s
a
b
Exptl.
Calc.
Slow step
3
3
3
3
3
a
00
05
10
15
20
3.21
3.83
4.52
5.17
6.05
3.40
4.55
5.25
5.06
5.80
6.66
4.67
3.55
2.89
3.41
Experimental
b
-2
-1
3
-1
Calculated:
k
obs were calculated by using k = 6.66 9 10
-1
s
,
K
1
= 1.02 dm ꢀmol
and
3
K
2
= 78.7 dm ꢀmol at 300 K in Eq. 7
Table 5 Activation parameters
for the oxidation of ketorolac by
hexacyanoferrate(III) in aqueous
alkaline medium at 300 K
Activation parameters
Values
overall reaction)
(
=
-1
DH (kJꢀmol
)
-1
29.9 ± 1
=
-1
DS (JꢀK ꢀmol
)
-220 ± 4
96 ± 2
=
-1
DG (kJꢀmol
)
log10
A
2.70 ± 0.04
operating in the present study as the reaction is first order in HCF and has fractional orders
-
in [KET] and [OH ]. Thus, fractional orders each in [KET] and [OH ] can be accom-
-
-
modated as in the scheme in which KET forms an anion by reacting with OH , and this
anion forms a complex with the oxidant in a pre equilibrium step, before the rate deter-
mining step where the complex decomposes to give a free radical.
The formation of free radicals has been assayed by the initial polymerization test. The
results show that test was negative in the absence of KET and positive in the presence of
KET; the overall reaction became strikingly faster when KET was present concurrently
-
with HCF, compared to KET alone in solution. The formation of a free radical from OH
is excluded, which is a quite natural process in alkaline KMnO oxidations [22]. Thus,
4
these results denote that HCF is capable of generating free radicals from KET, either
through its oxidation or, in a much slower way, through decarboxylation. The literature
reveals that HCF(III) oxidations of free radicals usually are a diffusion controlled step [23–
4
-
2
5]. Hence, HCF being a single equivalent oxidant, and its product [Fe(CN)6] did not
show any retarding effect, the generation of free radicals might be from the oxidative
decarboxylation of KET, as shown in Scheme 1.
The formation of complex between KET anion and HCF is proven kinetically by the
-
1
-1
non zero intercept graph of (kobs
)
against [KET]
(Michaelis–Menton plot). The
involvement of the ketonic oxygen and/or the nitrogen from the pyrrole ring in the adduct
formation with HCF can be ruled out, due to their steric hindrance. The participation of
carboxylic oxygen may be considered in adduct formation. However, it is known that HCF
-
is an inert complex and exchange of CN with the carboxylate anion of KET is quite
-
unusual due to non labile nature of CN . Nevertheless, HCF is reported to be [26] involved
in an outer sphere electron exchange with thio compounds through their intermediate
5
complex as [Fe(CN) –CNSO ] , particularly with substrates SO3 and S2O3 . Lancaster
-
2ꢁ
2ꢁ
5
3
123

924
J Solution Chem (2014) 43:916–929
O
C
O
N
N
C
K1
COO-
COOH
OH-
+
+
H O
2
(i)
O
C
N
K2
COO- + [Fe(CN)6]3-
Complex
(ii)
O
N
.
C
k
(
iii)
+ [Fe(CN) ]4- + CO2
Complex
6
Slow
O
C
O
C
N
N
.
+
fast
+
[Fe(CN) ]3-
+ [Fe(CN)6]4-
6
(iv)
O
C
O
C
N
N
+
fast
(v)
OH
OH-
+
Scheme 1 Mechanism of oxidation of ketorolac by hexacyanoferrate(III) in aqueous alkali
Table 6 Equilibrium constants involved in the mechanism of oxidation of ketorolac by hexacyanofer-
rate(III) in aqueous alkaline medium at different temperatures
Equilibrium
step
Equilibrium
constants
Temperature (K)
00 305
1.02 ± 0.02 1.21 ± 0.04 1.41 ± 0.04 1.60 ± 0.05 1.78 ± 0.05
79 ± 2 115 ± 3 139 ± 3 162 ± 4 183 ± 4
3
310
315
320
(Scheme 1)
i.
K
K
1
3
-1
-1
(
dm ꢀmol
)
)
ii.
2
3
(dm ꢀmol
and Murray [15] have suggested that many oxidations by transition metal complexes of
weak oxidants may proceed through intermediate complexes, as described in the case of
5
-
Fe(CN) –CNSO ] . In all such cases, the thio compounds are bonded with one CN
-
[
5 3
2
3
ꢁ
ligand through CN ꢁ SO : However, Leal et al. [27] proved that HCF decomposes to
-
CN with dissolved oxygen. Moreover, six ligands around the central metal atom do not
allow enough room for a seventh ligand in the octahedral complexes. Then, dissociative
processes are the most frequent in substitution reaction of octahedral complexes. Thus, in
1
23

J Solution Chem (2014) 43:916–929
925
Table 7 Thermodynamic parameters of various steps of the mechanism of oxidation of ketorolac by
hexacyanoferrate(III) in aqueous alkaline medium at 300 K
-
1
-1
-1
-1
Equilibrium step
DH* (kJꢀmol
)
DS* (JꢀK ꢀmol
)
DG* (kJꢀmol
)
i.
22.3 ± 1
74.5 ± 2
145 ± 3
-0.057 ± 0.004
-10.9 ± 0.5
ii.
32.5 ± 0.1
DH*, DS* and DG* refer to the concentration of the reactants maintained as in the reaction condition and for
temperature, 300 K)
-
the present study, the HCF may form a complex by exchanging CN with the oxygen of
carboxylate ion of KET, thereby the electron is transferred from KET to HCF giving a free
radical followed by decarboxylation. If that is the case, then the mechanism of oxidation is
of the inner sphere type. The small value of log10 A (2.70) clearly indicates that the reaction
is an inner sphere reaction. However, the values, ‘k’ calculated for the slow step are found
decrease with increasing temperature. This may be due to less easy formation of the KET
anion compared to the adduct at higher temperature (Table 4).
The rate law for the scheme can be derived as:
h
ꢁi
d½HCFꢂ
dt
3
6
ꢁ
rate ¼ ꢁ
However,
¼ k½complexꢂ ¼ kK1K2 FeðCNÞ
½KETꢂ ½OH ꢂ
ð2Þ
f
f
f
h
i
h
i
3
ꢁ
3ꢁ
6
FeðCNÞ6
¼ FeðCNÞ
þ½complexꢂ
T
f
h
i
h
i
3ꢁ
3ꢁ
ꢁ
¼
FeðCNÞ6 þK1K2 FeðCNÞ
½KETꢂ ½OH ꢂ
6
f
f
f
f
h
i
ꢀ
ꢁ
3
ꢁ
ꢁ
¼
FeðCNÞ6
1 þ K1K2½KETꢂ ½OH ꢂ
f
f
f
h
i
3
6
ꢁ
h
i
FeðCNÞ
3ꢁ
T
FeðCNÞ6
¼
ð3Þ
ð4Þ
ꢁ
f
f
1 þ K K ½KETꢂ ½OH ꢂ
1
2
f
ꢁ
ꢁ
ꢁ
½
¼
OH ꢂ ¼ ½OH ꢂ þ ½KET ꢂ
T
f
f
ꢁ
ꢁ
f
½OH ꢂ þ K1½OH ꢂ ½KETꢂ
f
f
ꢀ
ꢁ
ꢁ
¼
½OH ꢂ 1 þ K1½KETꢂ
f
f
ꢁ
½
OH ꢂ
ꢁ
f
½OH ꢂ ¼
f
1
þ K1½KETꢂ
f
-
At low concentration of KET, [OH ] = [OH ] , so that:
-
f
T
ꢁ
f
½
KETꢂ ¼ ½KETꢂ þ½KET ꢂ þ½complexꢂ
T
f
h
i
ꢁ
ꢁ
2ꢁ
ꢁ
¼
½KETꢂ þK1½OH ꢂ ½KET ꢂ þK1K2 FeðCNÞ
½OH ꢂ ½KETꢂ
f
f
f
6
f
f
f
ꢂ
h
i
ꢃ
ꢁ
2ꢁ
½OHꢁꢂ
¼
½KETꢂ þ 1 þ K1½OH ꢂ þK1K2 FeðCNÞ
f
f
6
f
f
½
KETꢂT
½
KETꢂ ¼ ꢂ
h
i
ꢃ
f
ꢁ
2ꢁ
½OHꢁꢂ
1
þ K1½OH ꢂ þK1K2 FeðCNÞ
f
6
f
f
123

926
J Solution Chem (2014) 43:916–929
3
-
In view of low concentration of [Fe(CN) ]
6
used in the experiments, the term
h
i
2
6
ꢁ
ꢁ
K1K2 FeðCNÞ
½OH ꢂ in the denominator of the right hand side can be neglected. Thus,
f
f
½
KETꢂT
þ K1½OH ꢂ
½
KETꢂ ¼ ꢂ
ꢃ
ð5Þ
ð6Þ
f
ꢁ
1
f
On substituting Eqs. 3, 4, 5 into Eq. 2:
h
i
3
ꢁ
ꢁ
T
kK1K2 FeðCNÞ6
½KETꢂ ½OH ꢂ
T
T
ꢁ
rate ¼ ꢀ
ꢁꢀ
ꢁ
ꢁ
f
1
þ K1K2½KETꢂ ½OH ꢂ 1 þ K1½OH ꢂ
f
f
ꢁ
kK1K2½KETꢂ ½OH ꢂ
T
T
kobs ¼
ꢁ
ꢁ
2
1
ꢁ 2
f
1
þ K1½OH ꢂ þK1K2½KETꢂ ½OH ꢂ þK K2½KETꢂ ½OH ꢂ
f
f
f
f
The rate law (Eq. 6) explains the first order dependence on [HCF] and the fractional
-
orders in [KET] and [OH ]
To verify the rate law, the approximation that the free and total concentrations are
substantially equal for Eq. 6 yields:
ꢁ
kK1K2½KETꢂ½OH ꢂ
kobs ¼
ð7Þ
ꢁ
ꢁ
2
1
ꢁ 2
1
þ K1½OH ꢂ þ K1K2½KETꢂ½OH ꢂ þ K K2½KETꢂ½OH ꢂ
and inversion of Eq. 7 gives:
ꢁ
1
1
1
1
K1½OH ꢂ
¼
þ
þ k
þ
ð8Þ
kobs
kK1K2½KETꢂ½OHꢁꢂ kK2½KETꢂ
The mechanism in Scheme 1 and corresponding rate law, in the form of Eq. 8, are verified
k
-1
- -1
-1
by plotting the graphs of (kobs versus [OH ] and [KET] at 300 K. They should be
linear and are found to be so (Fig. 2). From the slopes and intercepts of the plots, the values
)
-
2
-1
3
-1
of k, K and K are calculated as 6.66 (± 0.08) 9 10 ꢀs , 1.02 (± 0.02) dm ꢀmol and
1
2
3
-1
7
8.7 (± 0.3) dm ꢀmol , respectively, for 300 K.
Further, these values were used in the rate law (Eq. 7) under different experimental
conditions to regenerate kobs. The regenerated values are found to be in close agreement
with the experimental values (Table 1). This suppoprts the mechanism of oxidation shown
in Scheme 1.
It was observed that rate of reaction increases with increasing the ionic strength. But,
the slope (0.7) from the plot of log10 kobs against HI is less than the expected value of -4.0,
3
-
due to the interaction of triply charged HCF ([Fe(CN)6] ) and singly charged KET
carboxylate anion. This clearly indicates that the reaction follows an inner sphere mech-
anism through a common bridging ligand (via the carboxylate anion). This is further
-
2 -1
s .
supported by the relatively large value of ‘k’ of the slow step = 6.66 9 10
The study of the variation of relative permittivity of the reaction medium (D), varied by
varying the percentage compositions of t-butyl alcohol and water, on the rate of reaction
reveals that as the relative permittivity of the medium increases, the rate of reaction increases.
It can be concluded that the activated complex may be more solvated at higher dielectric
constant as the adduct is ionic in nature. Further, this activated complex may be more rigid
=
-1
-1
than the reactants, indicated by a large negative value of DS (-220 ± 4 JꢀK ꢀmol ) and a
= -1
large positive value of DG (96 ± 2 kJꢀmol ).
1
23

J Solution Chem (2014) 43:916–929
927
-
3
-1
1
/ [OH ] (dm · mol )
0
.0
1.0
2.0
3.0
4.0
1
1
5.0
2.0
7.0
6
5
4
3
2
1
.0
.0
.0
.0
.0
.0
9
6
3
0
.0
.0
.0
.0
0.0
12.0
0
.0
2.0
4.0
6.0
8.0
10.0
-
2
3
−1
1
/[KET] ×10 (dm ·mol )
Fig. 2 Verification of the rate law (Eq. 7) for the oxidation of ketorolac by hexacyanoferrate(III) in
aqueous alkaline medium at 300 K
The thermodynamic quantities for the equilibrium steps (i) and (ii) of the reaction scheme
-
can be evaluated as follows. The [KET] and [OH ] were varied at five different temperatures
-
1
in turn while keeping others constant. According to rate law (Eq. 8), the plots of (kobs)
-
against [OH ] and [KET] should be linear and were found to be so. From the slopes and
-1
-1
intercepts of such plots, the values of K were evaluated for all 5 temperatures (Table 6). A
1
-
1
van’t Hoff plot was made for the variation of K with temperature (log K against T ) and
1
10
1
the enthalpy of reaction (DH*) was evaluated from its slope. The DG* value was obtained
from the relation between DG* and equilibrium constant, K , for the temperature 300 K;
1
using DH* and DG*, the value of DS* was determined. Similarly, DH*, DS* and DG* (under
the reaction condition and temperature) were calculated for K for the equilibrium (ii)
2
(Table 7). The DH* and DS* values ofequilibria(i)and (ii) indicate that formation oftheKET
carboxylate anion and formation of a complex between oxidant and reductant is endothermic
in nature and they are formed by the loss of degrees of freedom. The small negative DG* value
of step (i) indicates that the formation of the KET carboxylate anion is kinetically governed
and the large value of DG* of step (ii) shows that formation of adduct is thermodynamically
-
3 -1
s
controlled, which is also evidenced by the small value of kobs (3.21 9 10
at 300 K).
4
Conclusion
Kinetics of oxidation of KET by HCF in alkali is routed through a formation of adduct,
formed between its anionic form with oxidant in a prior equilibrium step. This adduct is
decomposed by generating a free radical from KET. A small value of slope of log10 kobs
123

928
J Solution Chem (2014) 43:916–929
against HI plot indicates that the reaction mechanism is of the inner sphere type, which is
against the nature of HCF oxidations wherein it is reported as oxidations are usually of the
outer sphere type. The thermodynamic parameters indicate that the formation of the KET
anion and adduct are endothermic in nature.
Acknowledgments The authors are grateful to the Principal, Karnatak Science College, Dharwad, Kar-
nataka, India for providing the necessary facilities to carry out this work. They also thank the Raptakos Brett
and Co., Microlabs Ltd. KLAB, Mumbai, India for providing the free sample of Ketorolac.
References
1
2
3
. Kadurugamuwa, J.L., Hengstler, B., Zak, O.: Cerebrospinal fluid protein profile in experimental
pneumococcal meningitis and its attenuation by ampicillin and anti-inflammatory agents. J. Infect. Dis.
159, 26–34 (1989)
. Tuomanen, E., Hengstler, B., Rich, R., Bray, M.A., Zak, O., Tomasz, A.: Nonsterodial anti-inflam-
matory agents in the therapy for experimental pneumococcal meningitis. J. Infect. Dis. 155, 985–990
(1987)
. Guzman, A., Yuste, F., Toscano, R.A., Young, J.M., Van Horn, A.R., Muchowski, J.M.: Absolute
configuration of (-)-5-benzoyl-1,2-dihydro-3H-pyrrolo(1,2-a)pyrrole-l-carboxylic acid, the active
enantiomer of ketorolac. J. Med. Chem. 29, 589–591 (1986)
4
5
6
. Murry, M.D., Brater, D.C.: Renal toxicity of the nonsteroidal anti-inflammatory drugs. Annu. Rev.
Pharmacol. Toxicol. 33, 435–465 (1993)
. Alsarra, I.A., Bosela, A.A.: Proniosomes as a drug carrier for transdermal delivery of ketorolac. Eur.
J. Pharm. Biopharm. 59, 485–490 (2005)
. Kelson, E.P., Phengsy, P.P.: Kinetic study of 2-propanol and benzyl alcohol oxidation by alkaline
hexacyanoferrate(III) catalysed by a terpyridyl ruthenium complex. Int. J. Chem. Kinet. 32, 760–770
(2000)
7. Vovk, A.I., Muraveva, I.V., Kukhar, V.P., Baklan, V.F.: Kinetics of oxidation of vitamin B and its
1
O-acyl analogs with ferricyanide—A mechanistic model of thiamin-binding protein. Russ. J. Gen.
Chem. 70, 1108–1112 (2000)
8
9
. Tuwar, S.M., Nandibewoor, S.T., Raju, J.R.: Oxidation of chromium(III) by alkaline hexacyanofer-
rate(III). Trans. Met. Chem. 16, 335–338 (1991)
. Jose, T.P., Nandibewoor, S.T., Tuwar, S.M.: Kinetics and mechanism of oxidation of vanillin by
hexacyanoferrate(III) in aqueous alkaline medium. J. Solution Chem. 35, 51–62 (2006)
1
0. Billalli, H.B., Sharanabasamma, K., Tuwar, S.M.: Mechanism of the osmium(VIII)-catalysed oxidation
of fursemide by alkaline hexacyanoferrate(III) and analysis of fursemide by a kinetic and catalytic
method. Prog. React. Kin. Mech. 4, 347–367 (2010)
11. Leal, J.M., Garcia, B., Domingo, P.L.: Outer-sphere hexacyanoferrate(III) oxidation of organic sub-
strates. Coord. Chem. Rev. 173, 79–131 (1998)
1
2. Jose, T.P., Angadi, M.A., Salunke, M.S., Tuwar, S.M.: Oxidative study of Gabapentin by alkaline
hexacyanoferrate(III) in room temperature in presence of catalytic amount of ruthenium(III)—A
mechanistic approach. J. Mol. Struct. 892, 121–124 (2008)
1
1
3. Gutmann, V., Gritzner, G., Danksagmuller, K.: Solvent effects on the redox potential of hexacyano-
ferrate(III)–hexacyanoferrate(II). Inorg. Chim. Acta 17, 81–86 (1976)
4. Leal, J.M., Domingo, P.L., Garcia, B., Ibeas, S.: Alkali metal ion catalysis of the oxidation of L-ascorbic
acid by hexacyanoferrate(III) in strongly acidic media. J. Chem. Soc. Faraday Trans. 8, 3571–3577
(1993)
1
1
5. Lancaster, J.M., Murry, R.S.: The ferricyanide–sulphite reaction, J. Chem. Soc. A 2755–2758 (1971)
III
3 5
6. Martinez, M., Pitarque, M., Eldik, R.V.: Outer sphere redox reactions of [Co (NH )
(m-3)-
(H
x
y
P O
z
)]
complexes. A temperature and pressure-dependance kinetic study on the influence of
the phosphorus oxoanions. J. Chem. Soc. Dalton Trans. 2665–2671 (1996)
1
7. Khoshtariya, D.E., Kjaer, A.M., Marsagishvilli, T.A., Ulstrup, J.: Inhibition of near-infrared photo
induced electron transfer between free hexacyanoferrate(II) and hexacyanoferrate(III) by tetramethyl-
ammonium ions. J. Phys. Chem. 96, 4154–4156 (1992)
18. Jeffery, G.H., Bassett, J., Mendham, J., Denney, R.C.: Vogel’s Text Book Of Quantitative Chemical
Analysis, 5th edn. ELBS Longman, Essex, UK, (1991), (a) pp. 181 and 320
1
23

J Solution Chem (2014) 43:916–929
929
1
2
2
2
9. Feigel, F., Anger, V.: Spot Tests in Organic Analysis, translated by Oesper. R.E. Elsevier India Private
Limited, New Delhi (2005)
0. Hugar, G.H., Nandibewoor, S.T.: Kinetics of osmium(VIII) catalysis of periodate-oxidation of DMF in
aqueous alkaline medium. Transit. Met. Chem. 19, 215–217 (1994)
1. Garcia, B., Ruiz, R., Leal, J.M.: Kinetic study of the hexacyanoferrate(III) oxidation of dihydroxy-
fumaric acid in acid media. J. Phys. Chem. 112, 4921–4928 (2008)
2. Bilehal, D.C., Kulkarni, R.M., Nandibewoor, S.T.: Kinetic and mechanistic study of ruthenium(III)
catalysed oxidative deamination and decarboxylation of L-valine by alkaline permanganate. Can.
J. Chem. 79, 1926–1933 (2001)
(
2-)
23. Steenken, S., Neta, P.: Oxidation of substituted alkyl radicals by hexachloroiridate , hexacyanofer-
(3-)
rate , and permanganate ions in aqueous solution. Electron transfer versus chlorine transfer from
hexachloroiridate(2-) ion. J. Am. Chem. Soc. 104, 1244–1248 (1982)
2
2
2
2
4. Steenken, S., Buschek, J., MeClelland, R.A.: Radiation-chemical production and lifetimes of trialk-
oxymethyl carbocations in aqueous solution. J. Am. Chem. Soc. 108, 2808–2813 (1986)
5. Hughes, G., Willis, C.: Scavenger study in the radiolysis of aqueous ferricyanide solutions at high pH.
Discuss. Faraday Soc. 36, 223–231 (1963)
6. Swinehart, J.H.: The kinetics of the hexacyanoferrta(III) sulphite reaction. J. Inorg. Nucl. Chem. 29,
2313–2320 (1967)
7. Domingo, P.L., Garcia, B., Leal, J.M.: Acid-base behavior of the ferricyanide ion in perchloric acid
media. Spectrophotometric and kinetic study. Can. J. Chem. 68, 228–235 (1990)
123