Graphical separation of un-catalyzed and catalyzed reactions in iridium(III) catalyzed oxidation of cinnamaldehyde by cerium(IV) in aqueous acidic medium

Article abstract of DOI:10.1016/j.molcata.2008.07.007

When used as a homogeneous catalyst, iridium trichloride was found to change order of the reaction with respect to cerium(IV) sulphate concentrations in the oxidation of 3-phenyl propanal in aqueous acidic medium. Cerium(IV) and organic substrate form complex in the first equilibrium step. This complex decomposes in the rate determining step when iridium(III) is not present in the system while in the presence of iridium(III), it gives rise to another complex, which ultimately gives rise to cinnamic acid as product of reaction. Order of the reaction in the absence of iridium(III) was found to be one with respect to cerium(IV) sulphate concentrations while the reaction follows first order kinetics at low concentrations becoming to zero order at higher concentrations of the oxidant in the presence of catalyst. Both in the presence and the absence of catalyst, after showing first order kinetics in the beginning reaction tends to become independent of organic substrate at its higher concentrations. Rate of the reaction follows direct proportionality with respect to catalyst concentrations. Change in the ionic strength of the medium or concentrations of hydrogen ions, cerium(III) and acetic acid have no effect on the rate. Interestingly with fair degree of accuracy, the rate of un-catalyzed path may be separated graphically with the help of the intercept of catalyst graph. IR spectrum of the product confirmed cinnamic acid as the final oxidation product. Energy of activation, free energy of activation and entropy parameters were calculated.

Full text of DOI:10.1016/j.molcata.2008.07.007

Journal of Molecular Catalysis A: Chemical 293 (2008) 39–44
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Journal of Molecular Catalysis A: Chemical
journal homepage: www.elsevier.com/locate/molcata
Graphical separation of un-catalyzed and catalyzed reactions in iridium(III)
catalyzed oxidation of cinnamaldehyde by cerium(IV) in aqueous
acidic medium
Praveen K. Tandon∗, Manisha Purwar, Satpal Singh, Nidhi Srivastava, Manish Srivastava
Department of Chemistry, University of Allahabad, Allahabad 211002, India
a r t i c l e i n f o
a b s t r a c t
Article history:
When used as a homogeneous catalyst, iridium trichloride was found to change order of the reaction
with respect to cerium(IV) sulphate concentrations in the oxidation of 3-phenyl propanal in aqueous
acidic medium. Cerium(IV) and organic substrate form complex in the first equilibrium step. This complex
decomposes in the rate determining step when iridium(III) is not present in the system while in the
presence of iridium(III), it gives rise to another complex, which ultimately gives rise to cinnamic acid as
product of reaction. Order of the reaction in the absence of iridium(III) was found to be one with respect to
cerium(IV) sulphate concentrations while the reaction follows first order kinetics at low concentrations
becoming to zero order at higher concentrations of the oxidant in the presence of catalyst. Both in the
presence and the absence of catalyst, after showing first order kinetics in the beginning reaction tends to
become independent of organic substrate at its higher concentrations. Rate of the reaction follows direct
proportionality with respect to catalyst concentrations. Change in the ionic strength of the medium or
concentrations of hydrogen ions, cerium(III) and acetic acid have no effect on the rate. Interestingly with
fair degree of accuracy, the rate of un-catalyzed path may be separated graphically with the help of the
intercept of catalyst graph. IR spectrum of the product confirmed cinnamic acid as the final oxidation
product. Energy of activation, free energy of activation and entropy parameters were calculated.
© 2008 Elsevier B.V. All rights reserved.
Received 20 June 2008
Accepted 15 July 2008
Available online 25 July 2008
Keywords:
Cinnamaldehyde
Cerium(IV) sulphate
Iridium(III) chloride
Catalysis
1. Introduction
In continuation of these studies we hereby report the effect of irid-
ium(III) chloride in catalyzing the oxidation of cinnamaldehyde
Iridium(III) chloride does not show efficient catalytic activity in
alkaline medium [1,2]. Cerium(IV) has been reported to oxidize var-
ious aromatic compounds having different functional groups, but
these oxidations have been studied mainly from the synthetic point
of view [3]. Redox reactions of aromatic compounds are very impor-
tant from the synthetic point of view because most of the industries
are based on the oxidation products of various aromatics. Kinetic
study of the oxidation of various aromatic compounds helps to find
out the optimum conditions or concentrations under which maxi-
mum yield of the product may be obtained under most economical
conditions. Powerful catalytic activity of iridium(III)–cerium(IV)
system in acidic medium has been reported in the oxidation of
aliphatic ketones and alcohols [4–8]. This system oxidizes various
aromatic organic compounds in excellent yields, including cyclo-
hexane and benzene with the highest yields reported so far [9,10].
by cerium(IV) sulphate in aqueous sulphuric acid medium in the
presence and absence of iridium(III) chloride.
2. Experimental
Cerium(IV) sulphate, cinnamaldehyde (Lancaster), sulphuric
acid, acetic acid, ferroin (E. Merck) and cerium(III) sulphate (Fluka
A.G.) were used as supplied without further purification by prepar-
ing their solutions in double distilled water. Strength of sodium
hexachloroiridate(III) (Johonson Matthay & Co.) was 3.35 × 10−3 M
which was prepared by dissolving the sample in minimum amount
of AR HCl (0.62 × 10−2 M). Ferrous ammonium sulphate (FAS) solu-
tion was standardized by titrating it against a standard solution
of potassium dichromate (E. Merck) using N-phenyl anthranilic
acid as an internal indicator. Stock solution of cerium(IV) sul-
phate was prepared by dissolving the sample in AR sulphuric
acid (diluted with distilled water in 1:1 ratio) and its strength
was checked by titrating it against a standard ferrous ammonium
sulphate solution using ferroin as an internal indicator. Strength
of sulphuric acid present in cerium(IV) sulphate was adjusted in
∗
Corresponding author. Tel.: +91 9415310943; fax: +91 532 2461236.
E-mail addresses: pktandon123@rediffmail.com,
ptandonk@yahoo.co.in (P.K. Tandon).
1381-1169/$ – see front matter © 2008 Elsevier B.V. All rights reserved.
doi:10.1016/j.molcata.2008.07.007

40
P.K. Tandon et al. / Journal of Molecular Catalysis A: Chemical 293 (2008) 39–44
2.1. Determination of kinetic orders
Oxidation of cinnamaldehyde by cerium(IV) sulphate in acidic
medium takes place even in the absence of iridium(III) chloride.
Therefore, un-catalyzed and catalyzed reactions were studied sep-
arately under similar conditions. In the case of oxidant variation
−dc/dt values were calculated at a fixed initial time from the indi-
vidual plots, while in all other cases values were calculated at a fixed
initial concentration. Rate values (−dc/dt), obtained from the initial
slopes of individual graphs between the residual concentrations of
cerium(IV) at various time intervals, were finally plotted against the
changing concentrations of the particular reactant for which order
of the reaction was to be obtained. Orders, with respect to various
reactants were confirmed by plotting log of remaining concentra-
tion of oxidant versus time (oxidant variation), by plotting −dc/dt
values versus concentration of the reactant, by calculating slope of
the double logarithmic graphs between rate versus concentration
and by calculating rate constant for molar concentrations (catalyst
variation). Study could not be made at constant ionic strength of
the medium due to large volumes of potassium chloride required
to keep the ionic strength constant. However, effect of ꢀ on the
rate was studied separately with the help of a standard solution of
potassium chloride.
Fig. 1. Sample individual time plots for the consumption of CeIV at 30 ◦C.
[H2SO4] = 0.15 M,
[cinnamaldehyde] = 1.00 × 10−3 M,
[acetic
acid] = 0.09 M,
2.2. Product study and stoichiometry
[IrCl3] = 0.67 × 10−6
M (for catalyzed reactions only) and [Ce(SO4)2] = A and
C—0.40; B and D—1.00 (×10−3 M). (Concentrations are the initial concentration of
reactants.)
To find out stoichiometry, reactions were studied under the
conditions when cerium(IV) sulphate was in large excess com-
pared to organic substrate. Reaction mixtures were kept at 30 ◦C
for a long period to ensure complete oxidation of the organic
substrate. After completion of the reaction, solution was concen-
trated on a water bath. Reaction mixture was then extracted with
diethyl ether (5× 25 ml) and the solvent evaporated under reduced
pressure. After re-crystallization with hot water compound was
obtained as white solid (mp 130 ◦C, reported 133 ◦C), which gave
the test for unsaturation [11]. The compound was finally identified
by chromatographic technique [12] and also by taking IR spectra
of the product. Peaks at 3027 cm−1 (ꢁO–H); 1680 cm−1 (ꢁC O);
1494 cm−1 (ꢁC C) in the IR spectrum (ABB Bomen, FTLA 2000 104)
of the product confirmed the formation of cinnamic acid in the
reaction mixture. Lowering of the peak (observed at 3027 cm−1) of
hydroxyl group in acids is mainly due to overlap of C–H stretch.
[13] Stoichiometry of the reaction may be given by the following
equation:
maintaining sulphuric acid concentration in the reaction mixture.
All other chemicals used were of either AR grade or chemi-
cally pure substances. Cinnamaldehyde was insoluble in aqueous
medium therefore its stock solution was prepared by dissolving
the sample in minimum amount of acetic acid and then dilut-
ing the solution to the desired volume. Concentration of acetic
acid was kept constant in all the variations except while study-
ing the effect of acetic acid concentration on the rate. Progress
of the reaction was measured (constant temperature 0.1 ◦C) at
different intervals of time by transferring 5.0 ml of aliquot of
reaction mixture to the fixed amount of ferrous ammonium sul-
phate solution (volume or strength of FAS solution should be
slightly in excess to the total cerium(IV) initially taken in the
reaction mixture) and estimating the excess ferrous ammonium
sulphate with a standard cerium(IV) sulphate solution using fer-
roin as an internal indicator. Thus, titre values directly correspond
to the amount of cerium(IV) consumed in the reaction mixture.
In all kinetic runs cinnamaldehyde concentration was kept in
excess.
C6H5CH CHCHO + 2CeIV + H2O
→ C6H5CH CHCOOH + 2CeIII + 2H+
(1)
Table 1
Effect of variation of [cerium(IV)] and [cinnamaldehyde] on the rate at 30 ◦C
#
[CeIV
]
(×103 M)
−dc/dt (×105 M min−1
)
k((−dc/dt)/[CeIV])
[Cinnamaldehyde]* −dc/dt (×105 M min−1
(×103 M)
)
k((−dc/dt)/[cinnamaldehyde])
(×102 min−1
)
(×103 min−1
)
(A)
(B)
(A)
(B)
(A)
(B)
(A)
(B)
0.40
0.50
0.60
0.70
0.80
0.90
1.00
–
0.54
0.62
0.63
0.65
0.68
0.71
0.71
–
0.10
0.12
0.14
0.18
0.20
0.23
0.24
–
1.35
1.24
1.05
0.92
0.85
0.78
0.70
–
0.25
0.24
0.23
0.25
0.25
0.25
0.24
–
1.00
1.25
1.50
2.00
2.50
3.30
4.00
5.00
0.50
0.58
0.63
0.63
0.72
0.82
0.88
1.14
0.34
0.36
0.38
0.42
0.50
0.67
0.80
0.88
5.00
4.64
4.20
3.45
2.88
2.48
2.20
2.28
3.40
2.88
2.53
2.10
2.00
2.00
2.00
1.76
[Ce(SO4)2] = 1.00 × 10−3 M (for * only), [cinnamaldehyde] = 1.0 × 10−3 M (for # only). [H2SO4] = 0.15 M, [IrCl3] = 0.67 × 10−6 M (for A only), [acetic acid] = 0.50 M (for *); 0.09 M
(for #). (Concentrations are the initial concentration of reactants.)

P.K. Tandon et al. / Journal of Molecular Catalysis A: Chemical 293 (2008) 39–44
41
Fig. 4. Effect of variation of [catalyst] on the rate at 30 ◦C. [Ce(SO4)2] = 0.67 × 10−3 M,
[cinnamaldehyde] = 2.00 × 10−3 M, [H2SO4] = 0.15 M and [acetic acid] = 0.15 M. (Con-
centrations are the initial concentration of reactants.)
Fig. 2. Effect of variation of [cerium(IV)] on the rate at 30 ◦C. [H2SO4] = 0.15 M,
lows first order kinetics in the case of un-catalyzed reaction while
in the case of catalyzed reaction, the rate shows direct propor-
tionality with respect to oxidant only at its low concentrations
and tends to become independent of concentration at higher con-
centrations of the oxidant. In Table 1, −dc/dt values increase with
increasing concentration of organic substrate in the beginning but
at higher concentrations the increase is not so prominent. This
trend becomes clear on plotting −dc/dt values versus [ketone] in
which line appears to be passing through the origin in the begin-
ning at low concentrations. At higher concentrations of organic
substrates the line tends to become parallel to the x-axis (Fig. 3).
These facts indicate that in both cases, order of the reaction with
respect to organic substrate is one at low concentrations but tends
to become zero at higher concentrations of the organic substrate.
Study at still lower concentrations of organic substrate could not be
performed due to practical difficulties as the reaction becomes too
slow to be measured properly, moreover, lowering the concentra-
tions further, may not be in accordance with the Ostwald’s isolation
method, yet the trend of lines is quite clear from Fig. 3.
[cinnamaldehyde] = 1.00 × 10−3 M, [acetic acid] = 0.09 M and [IrCl3] = 0.67 × 10−6
M
(for catalyzed reactions only). (Concentrations are the initial concentration of
reactants.)
3. Results
First order plots between the log of residual oxidant concen-
tration in the reaction mixture versus time in case of uncatalyzed
reactions show parallel straight lines, however, probably due to the
formation of complex, in the case of catalyzed reaction, deviations
in the later part of the reaction at higher concentrations of the
oxidant are more pronounced (Fig. 1). In the case of un-catalyzed
reaction −dc/dt values increase proportionately with increasing
concentration of the oxidant, while in the case of catalyzed reaction
−dc/dt values increase proportionately with concentration only in
the beginning and the increase becomes less prominent at higher
concentrations (Table 1). First order rate constant values for molar
concentrations of the oxidant show fair constancy in the case of
un-catalyzed reaction, while these values remain constant only in
the beginning in the case of catalyzed reaction. On plotting −dc/dt
values versus concentration of the oxidant straight lines passing
through the origin are obtained in the case of both catalyzed and
un-catalyzed reactions. In the case of un-catalyzed reaction devi-
ation from linearity in the graph was not observed even at higher
concentrations of the oxidant, while in the case of catalyzed reac-
tion the straight line passing through the origin becomes parallel
to x-axis at higher concentrations of the oxidant (Fig. 2). All these
facts collectively confirm that rate of the reaction through out fol-
On plotting −dc/dt values versus concentration of [IrCl3] (Fig. 4),
a straight line was obtained with a positive intercept at y-axis. Slope
of this line shows the rate of catalyzed reaction while the positive
intercept gives the extent of the reaction, which takes place even in
the absence of the catalyst. Proportionate increase in −dc/dt values
with increasing catalyst concentrations and fair constancy in the
first order rate constant k values indicate that the reaction follows
first order kinetics with respect to iridium(III) chloride concentra-
tions (Table 2). On plotting double logarithmic graphs between
log −dc/dt and log[IrCl3] straight lines with slope value 0.95 was
obtained, which further supports direct proportionality of the reac-
tion velocity with respect to iridium(III) chloride concentrations.
−dc/dt values remain practically constant at 0.59 0.03
and 0.36 0.02 (×10−5 M min−1
) in the case of catalyzed
Table 2
Effect of variation of [iridium(III)] on the rate at 30 ◦C
[IrIII] (×106 M)
−dc/dt (×105 M min−1
)
k((−dc/dt)/[IrIII]) (min−1
)
0.34
0.67
1.00
1.34
1.50
1.67
2.00
0.29
0.35
0.40
0.45
0.47
0.50
0.56
8.50
5.20
4.00
3.35
3.13
2.94
2.73
Fig. 3. Effect of variation of [cinnamaldehyde] on the rate at 30 ◦C. [Ce(SO4)2] =
1.00 × 10−3 M, [H2SO4] = 0.15 M, [acetic acid] = 0.50 M and [IrCl3] = 0.67 × 10−6 M (for
catalyzed reactions only). (Concentrations are the initial concentration of reactants.)
[Ce(SO4)2] = 0.67 × 10−3 M, [cinnamaldehyde] = 2.00 × 10−3 M, [H2SO4] = 0.15 M,
[acetic acid] = 0.15 M. (Concentrations are the initial concentration of reactants.)

42
P.K. Tandon et al. / Journal of Molecular Catalysis A: Chemical 293 (2008) 39–44
and un-catalyzed reactions respectively on changing sulphuric
acid concentrations from 0.20 to 0.75 M (under the condi-
tions [IrCl3] = 0.67 × 10−6 M (in case of catalyzed reaction only)
[Ce(SO4)2] = 1.00 × 10−3 M; [cinnamaldehyde] = 2.0 × 10−3 M and
[acetic acid] = 0.15 M). On changing the concentration of exter-
nally added Ce2(SO4)3 from 1.00 to 2.00 × 10−3 M the rate values
ferent thermodynamic parameters were calculated. The sequence
in the rate of oxidation is also confirmed from the energy of acti-
vation, entropy of activation and free energy of activation values,
which were found to be 12.6 and 8.64 (K cal), −48.00 and −57.00
(e.u.), 17.28 and 14.55 (kcal g−1 mole−1) for un-catalyzed and cat-
alyzed reactions, respectively.
remain constant at 0.72 0.02 and 0.41 0.01 (×10−5 M min−1
)
in the case of catalyzed and un-catalyzed reactions respec-
tively (under the conditions [IrCl3] = 0.67 × 10−6 M (in case of
catalyzed reaction only), [Ce(SO4)2] = 1.00 × 10−3 M; [cinnamalde-
hyde] = 2.0 × 10−3 M, [H2SO4] = 0.25 M and [acetic acid] = 0.15 M).
Constancy in the rate values shows that changes in the concen-
tration of hydrogen ions or Ce(III) ions do not effect the reaction
velocity. These observations show that these ions are not given
out before the rate determining slow step. Similarly on changing
the concentration of acetic acid from 0.15 to 0.75 M (in case of un-
catalyzed reaction 0.25 to 1.00 M), the rate values remain constant
at 0.58 0.028 and 0.33 0.019 (×10−5 M min−1) in the case of
catalyzed and un-catalyzed reactions respectively (under the con-
ditions [IrCl3] = 0.67 × 10−6 M (in case of catalyzed reaction only),
[Ce(SO4)2] = 1.00 × 10−3 M; [cinnamaldehyde] = 2.00 × 10−3 M and
[H2SO4] = 0.15 M). Large volumes of potassium chloride required
to maintain ionic strength of the medium constant, restricted
the study to be conducted at constant ionic strength of the
medium. However, effect of change of ꢀ on the reaction rate
was studied separately. Change in ionic strength of the medium
with the help of a standard solution of potassium chloride does
not affect the reaction velocity and the rate values (−dc/dt)
4. Reactive species of iridium(III) chloride
3−
It is known that IrCl3 in hydrochloric acid gives IrCl6 species
[14]. It has also been reported that iridium(III) and iridium(I)
ions are the stable species of iridium [15]. Further, the aquation
of [IrCl63−] gives [IrCl5(H2O)2−], [IrCl4(H2O)2]− and [IrCl3(H2O)3]
species [16–18]. This equilibrium may be shown in general by the
following equation:
IrCl63− + n-H2O ꢀ [IrCl6−n(H2O)n]3−n + Cl−
(2)
No effect of chloride ions on the reaction rate in the present study
indicates that the above equilibrium is shifted more towards the left
side and IrCl5(H2O)2− cannot be considered as the reactive species
3−
[1,2]. Therefore, considering our experimental results, IrCl6 has
been considered to be the reactive species of iridium(III) chloride
in the present study, which has been considered to be the active
species previously also [1,2,6]. Further, change in the oxidation state
of iridium during the course of the reaction may also result in the
removal of chloride ion(s), leading to their negative effect on the
reaction velocity, which was not observed in the present case indi-
cating that change in the oxidation state of iridium also may not be
possible.
remain constant at 0.58 0.017 and 0.40 0.02 (×10−5 M min−1
)
in the case of catalyzed and un-catalyzed reactions respec-
tively (under the conditions [IrCl3] = 0.67 × 10−6 M (in case of
catalyzed reaction only), [Ce(SO4)2] = 1.00 × 10−3 M; [cinnamalde-
hyde] = 2.00 × 10−3 M, [H2SO4] = 0.15 M, [acetic acid] = 0.15 M). This
observation shows that changes in the ionic strength of the medium
and the concentration of chloride ions have no effect on the reaction
velocity. Arrhenius equation was found to be applicable and from
the slopes of the Arrhenius plots and by using Eyring equation, dif-
5. Discussion
Formation of 1:1 complexes between cerium(IV) and alcohols
[19] and ketones [20] are well documented. Complex formation
between cerium(IV) and alcohols to give complexes of composi-
tion [ROH·cerium(IV)]4+ and Michaelis–Menten type of kinetics has
Scheme 1. Proposed steps in the oxidation of cinnamaldehyde by cerium(IV) in aqueous acidic medium in the presence of iridium(III) chloride.

P.K. Tandon et al. / Journal of Molecular Catalysis A: Chemical 293 (2008) 39–44
43
been reported [21]. While similar results have been reported in the
case of ketones also [22–24].
Rate in terms of decreasing concentration of cerium(IV) from step
(III) of the mechanism may be given as
d[CeIV
dt
]
kK1K2[CeIV][S][IrIII
1 + K1K2[CeIV][S]
]
−
=
(5)
5.1. Mechanism in the presence of iridium(III) chloride
This equation explains all experimental findings. At low con-
centrations of oxidant and organic substrate the inequality
1 ꢀ K1K2[CeIV][S] may hold and Eq. (5) reduces to (6) which,
explains the nature shown by various reactants.
Thus, according to Scheme 1, cerium(IV) species combines with
organic substrate to give complex C1, which in turn combines with
iridium(III) species to give the complex C2. Complex C2 in the slow
and rate determining step gives rise to the intermediate product,
which ultimately converts into p-methoxy benzoic acid. Electron
of the hydrogen atom is quickly taken up by another cerium(IV),
giving rise to the original iridium(III) species.
d[CeIV
]
−
= kK1K2[CeIV][S][IrIII
]
(6)
dt
At higher concentrations of oxidant and substrate the reverse
inequality 1 ꢁ kK1K2[CeIV][S] holds good and the equation becomes
5.2. Mechanism in the absence of iridium(III) chloride
as
d[CeIV
dt
]
In the case of un-catalyzed path all the results were similar
except of the order with respect to cerium(IV) concentrations, in
which the reaction shows direct proportionality with respect to
oxidant even at higher concentrations of the oxidant. Thus on the
basis of similar results reaction path in the absence of catalyst may
be given as according to Scheme 2.
−
= k[Ir(III)]T
(7)
Further verification of the rate law (5) may be given by rewriting
Eq. (8) as
1
1
1
=
+
(8)
−d[CeIV]/dt
kK1K2[CeIV][S]
k[IrIII
]
Values of kK1K2 were calculated from the slopes of the graphs
between 1/rate versus 1/[CeIV], 1/[substrate] and 1/[IrIII], which
comes out to be 3.15, 1.49 and 1.67 (×107) respectively. Fair con-
stancy in the kK1K2 values calculated from three different graphs
further indicates the validity of the Scheme 1 and the rate law (5).
It may be pointed out that the intercept of the catalyst graph gives
the extent of the reaction, which takes place even in the absence of
iridium(III) chloride. This value was calculated and was found to be
1.64 for the un-catalyzed reaction.
6. Rate laws
6.1. Derivation of rate law for the catalyzed path
Considering the equilibrium concentrations of complexes C1
and C2 in steps (I) and (II) of the mechanism 1 and putting concen-
tration of C1 from step (I) into the concentration of [IrIII] obtained
from step (II), total concentration of catalyst may be given as accord-
ing to the following equation:
[C2]
[IrIII
]
=
+ [C2]
]
(3)
6.2. Derivation of rate law for un-catalyzed path
Total
K1K2[S][CeIV
From Eq. (3) concentration of complex C2 is given as
K1K2[CeIV][S][IrIII
Similarly for the un-catalyzed path total concentration of
cerium(IV) may be given as
[C1]
K1ꢂ [S]
]
[CeIV] =
(9)
T
[C2] =
(4)
1 + K1K2[CeIV][S]
Scheme 2. Proposed steps in the oxidation of cinnamaldehyde by cerium(IV) in aqueous acidic medium in the absence of iridium(III) chloride.

44
P.K. Tandon et al. / Journal of Molecular Catalysis A: Chemical 293 (2008) 39–44
and considering equilibrium in step 1 of the mechanism 2 concen-
tration of cerium(IV) may be given as
medium when used in conjunction with cerium(IV) sulphate. The
present iridium(III)–cerium(IV) system is especially useful in the
synthetic reaction of industrial importance. The present study also
proves that the rate and extent of un-catalyzed reaction can be cal-
culated with fair degree of accuracy by the intercept of the catalyst
graph by studying the catalyzed reaction only.
[CeIV
]
= [CeIV] + [C1]
(10)
Total
Putting the value of [Ce(IV)] from Eq. (9) into Eq. (10) [C1] comes
out to be
K1ꢂ [S][CeIV
]
T
[C1] =
(11)
1 + K1ꢂ [S]
Acknowledgement
and the rate in terms of decreasing concentrations of cerium(IV)
sulphate may be given as
M.S. gratefully acknowledges the financial assistance by the Uni-
versity Grants Commission, New Delhi (grant no. F-32-240/2006
(SR).
kꢂK1ꢂ [S][CeIV
1 + K1ꢂ [S]
]
T
d[CeIV
dt
]
−
= kꢂ[C1] =
(12)
Appendix A. Supplementary data
This equation explains first order kinetics with respect to organic
substrate, which tends to become zero order at higher concentra-
tions of organic substrate, and the first order kinetics with respect to
cerium(IV) concentrations even at its higher concentrations. Devi-
ations from the straight line at higher concentrations of oxidant
as observed in Fig. 2 are also explained due to the formation of
complex between the oxidant and organic substrate in the first
step of the mechanism 2. Rate law given by Eq. (13) reduces to
as given in Eq. (14) if the inequality 1 ꢀ K1[S] is considered at low
concentrations of organic substrate.
Supplementary data associated with this article can be found,
in the online version, at doi:10.1016/j.molcata.2008.07.007.
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d[CeIV
dt
]
−
= kꢂK1ꢂ [S][CeIV
]
(13)
T
Eq. (13) may also be written in the form
1
rate
1
1
=
+
(14)
kꢂK1ꢂ [CeIV][S]
kꢂ[CeIV
]
On plotting graphs between 1/rate versus 1/[S] or 1/[CeIV]T with the
help of Eq. (14) we should get straight lines with positive intercepts
at y-axis. The values of kꢂK1ꢂ were calculated from two graphs and
were found to be 2.20 and 1.12, while this value calculated from
the intercept of the catalyst graph in the catalyzed reaction was
found to be 1.64. Close similarity in the rate values calculated from
two separately studied reactions finally confirm the validity of the
mechanisms and the rate laws proposed for both the reactions. The
study also provides a method by which with fair degree of accuracy,
the rate and extent of the un-catalyzed reaction can be calculated
by studying the catalyzed reaction only.
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Absence of any reaction between the organic substrate and irid-
ium(III) giving rise to the complex C1 in first step of Scheme 1 rules
out the possibility of interaction between these two. No effect of
chloride ions on the rate negates the possibility of their release
before the rate determining step. Thus we can safely assume the
validity of the final rate law and at least formation of the complexes
before the rate determining step.
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7. Conclusion
The present study shows that iridium(III) chloride acts as a very
effective catalyst for the oxidation of aromatic aldehydes in acidic