Oxidation of ketones by ceric perchlorate catalysed by iridium(III)

Article abstract of DOI:10.1016/j.molcata.2005.01.031

In search of economical and effective catalysts it was found that IrCl _3, a sluggish catalyst in alkaline media, surpasses catalytic efficiency of osmium and ruthenium in aqueous perchloric acid medium in the oxidation of aliphatic ketones by ceric perchlorate. Rate decreases in the beginning at low acid concentrations, reaches to a minimum and then becomes proportional to [HClO₄], probably due to conversion of hydrolysed to unhydrolysed species of ceric perchlorate, which then accelerates the rate. Reduction of oxidant by water depends on the concentrations of acid and [cerium(IV)] initially used. Orders are one with catalyst, one initially with respect to oxidant and ketones tending to become zeroth order at higher concentrations. No effect of change of μ on the rate was observed. Thermodynamic data suggest that diethyl ketone (DET) forms the activated complex more easily compared to dimethyl ketone (DMT).

Full text of DOI:10.1016/j.molcata.2005.01.031

Journal of Molecular Catalysis A: Chemical 232 (2005) 83–88
Oxidation of ketones by ceric perchlorate catalysed by iridium(III)
Praveen K. Tandon∗, Sumita Sahgal, Alok K. Singh, Gayatri, Manisha Purwar
Department of Chemistry, University of Allahabad, Allahabad 211002, India
Received 12 November 2004; accepted 20 January 2005
Abstract
In search of economical and effective catalysts it was found that IrCl3, a sluggish catalyst in alkaline media, surpasses catalytic efficiency
of osmium and ruthenium in aqueous perchloric acid medium in the oxidation of aliphatic ketones by ceric perchlorate. Rate decreases in
the beginning at low acid concentrations, reaches to a minimum and then becomes proportional to [HClO4], probably due to conversion
of hydrolysed to unhydrolysed species of ceric perchlorate, which then accelerates the rate. Reduction of oxidant by water depends on the
concentrations of acid and [cerium(IV)] initially used. Orders are one with catalyst, one initially with respect to oxidant and ketones tending
to become zeroth order at higher concentrations. No effect of change of µ on the rate was observed. Thermodynamic data suggest that diethyl
ketone (DET) forms the activated complex more easily compared to dimethyl ketone (DMT).
© 2005 Elsevier B.V. All rights reserved.
Keywords: Iridium(III) chloride; Catalysis; Cerium(IV) perchlorate; Oxidation; Ketones
1. Introduction
alkaline medium [4]. It may be mentioned here that catalytic
activities of ruthenium(VIII) in acidic medium [5] and the
mechanistic steps involved in iridium(III) chloride catalysis
in alkaline medium [4] were reported for the first time from
our laboratory. Here we report the oxidation of dimethyl ke-
tone and diethyl ketone by cerium(IV) perchlorate in aqueous
perchloric acid medium catalysed by iridium(III) chloride.
Homogeneous catalysis by iridium(III) chloride in the ox-
idation of organic compounds has not been given attention
due to sluggish catalytic behaviour of iridium compounds in
alkaline medium. Interestingly it was observed by us that the
average concentration of iridium(III) chloride (10−7 M), re-
quired to catalyse the oxidation of similar ketones, is at least
hundred times less compared to the concentration of ruthe-
nium(III) chloride with cerium(IV) sulphate [1] and osmium
tetroxide with alkaline hexacyanoferrate(III) [2]. Formation
of hazardous osmates restricts the use of osmium as homoge-
neous catalyst to the alkaline medium only while, ruthenium
compounds get advantage as they can be used in acidic as
well as in alkaline medium both. Main difference between
the two lies in the fact that osmium adds to the double bond
while ruthenium compounds are reported to break the dou-
ble bond [3]. Homogeneous catalysis by iridium(III) chlo-
ride, belonging to the same group of periodic table, has been
given little attention due to its sluggish catalytic activity in
2. Experimental
Sodium perchlorate, cerium(IV) sulphate (Loba Chemie
Indaustranal Co.), sulphuric acid, ferrous ammonium
sulphate, ferroin, dimethyl ketone(DMT) (E. Merck) and
diethyl ketone (DET) (Baker Analysed) were used as such
without further purification and their solutions were pre-
pared by directly dissolving the weighed sample in doubly
distilled water. Solution of sodium hexachloroiridate(III)
(Johonson Matthay and Co.) was prepared by dissolving the
sample in the minimum amount of hydrochloric acid (A.R.);
the final strengths of acid and catalyst were 0.00624 and
3.35 × 10−3 M, respectively. The strength of cerium(IV)
sulphate, prepared by dissolving the sample in 1:1 sulphuric
∗
Corresponding author. Tel.: +91 532 2607435; fax: +91 532 2545021.
E-mail address: pktandon123@rediffmail.com (P.K. Tandon).
1381-1169/$ – see front matter © 2005 Elsevier B.V. All rights reserved.
doi:10.1016/j.molcata.2005.01.031

84
P.K. Tandon et al. / Journal of Molecular Catalysis A: Chemical 232 (2005) 83–88
acid, was determined by titrating it against a standard solution
of ferrous ammonium sulphate, using ferroin as an internal
indicator. Solution of ceric perchlorate was prepared by pre-
cipitating the ceric hydroxide [6] from the prepared solution
of ceric ammonium nitrate with dilute solution of ammonium
hydroxide. The gelatinous precipitate thus, obtained was
washed several times with distilled water till the complete
removal of sulphate ions and after drying the precipitate, it
was re-dissolved in perchloric acid. All other chemicals used
were either Analar or chemically pure substances. Progress of
the reaction was measured (at constant temperature 0.1 ◦C)
at different intervals of time by transferring the alliquot to
a fixed amount of ferrous ammonium sulphate solution (in
slight excess to cerium(IV) perchlorate initially taken) and
titrating the remaining ferrous ammonium sulphate, with a
standard cerium(IV) sulphate solution using ferroin as an
internal indicator. In this way the titre values directly corre-
spond to the amount of cerium(IV) perchlorate consumed in
the reaction mixture. In all kinetic runs ketone was present
in excess.
oxidation of the organic substrate was ensured. Total amount
of cerium(IV) perchlorate consumed by 1 mole of organic
substrate for its complete oxidation was determined. The
final products, identified with the help of TLC and the spot
test methods [7], were found to be the corresponding acids
in both the cases. The stoichiometry of the reaction may be
given according to the following equation
RCO + CeIV + H2O → RCOOH + CeIII + H+
(1)
where R = (CH3)2 and (C2H5)2 for DMT and DET,
respectively.
3. Results
Sample individual time plots between log(a − x) versus
time for the lowest and highest concentrations of [Ce(ClO4)4]
for two ketones show parallel straight lines (Fig. 1). In Table 1
the rate constant values obtained by two methods (kgraphical
and kcalculated) show constancy in the beginning, but the rate
values in both cases start decreasing gradually with increas-
ing [oxidant]. Trend in −dc/dt values shows that the values
increase proportionately with increasing [oxidant] in the
beginning but the increase is not prominent at higher concen-
trations. All these factors indicate that the reaction follows
first order kinetics at low concentrations, which tends to
become zeroth order at higher [oxidant]. This trend becomes
clearer on plotting −dc/dt values versus [Ce(ClO4)4] where
2.1. Determination of kinetic orders and stoichiometry
Iridium(III) chloride catalysed oxidation of dimethyl ke-
tone and diethyl ketone by cerium(IV) perchlorate in aqueous
perchloric acid was studied under the conditions and range
in which the uncatalysed reaction was negligible. Rate of
reaction (−dc/dt) was obtained from the initial slopes of in-
dividual graphs plotted between the residual concentration
of cerium(IV) at various time intervals. In case of oxidant
variation rate values were calculated at a fixed initial time
while in all other cases they were calculated at a fixed initial
concentration. The rate values (−dc/dt) thus obtained were
finally plotted against the changing concentrations of the par-
ticular reactant for which order of the reaction was to be
obtained. Close resemblance in the rate values obtained by
two methods, i.e. kcalculated (by using integrated first order
rate equation and taking average of values in a particular set)
and kgraphical (by dividing −dc/dt values with (a − x), i.e. the
residual concentration of cerium(IV) at the point at which
rates were calculated), confirms authenticity of the results.
Orders, with respect to various reactants were confirmed by
plotting log(a − x) versus time (oxidant variation), by plot-
ting −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). Tables and figures
contain initial concentrations of the reactants. Ionic strength
of the medium could not be kept constant in all the variations
due to large volumes of sodium perchlorate required to keep
the ionic strength constant. However, effect of µ on the rate
was studied separately with the help of a standard solution of
sodium perchlorate.
Fig. 1. Sample individual time plots for consumption of CeIV at 25 ◦C:
[HClO4] = 0.75 M, (A) [DMT] = 4.0 × 10-3 M, [IrCl3] = 5.0 × 10-7 M.
[Ce(ClO4)4] = A – 2.0; B – 12.0 (×10−4 M). (B) [DET] = 1.0 × 10-3 M,
[IrCl3] = 2.0 × 10-7 M, [Ce(ClO4)4] = C – 3.50; D – 20.0 (×10−4 M).
Stoichiometry of the reaction was studied by taking
cerium(IV) perchlorate in large excess compared to the
organic substrate in different ratios, and thus complete

P.K. Tandon et al. / Journal of Molecular Catalysis A: Chemical 232 (2005) 83–88
85
Table 1
Variation of [cerium(IV)] and [ketones] on the reaction rate at 25 ◦C
[CeIV]# ×
−dc/dt ×
kgr.
×
kcal.
×
[Ketone]*
103
×
−dc/dt ×
kgr.
×
kcal.
×
104 M min−1
106 M min−1
102 M min−1
102 M min−1
106 M min−1
102 min−1
102 min−1
A
B
A
Bꢀ
A
B
A
B
A
B
A
B
2.0
3.0
3.5
4.0
4.5
5.0
5.5
7.0
7.5
2.22
2.92
3.46
3.85
4.50
5.00
–
5.00
–
5.00
–
5.25
5.00
–
–
–
–
–
1.27
1.08
1.08
1.07
1.10
1.10
–
0.81
–
0.59
–
0.50
0.45
–
–
–
–
–
1.22
1.00
1.03
1.10
1.15
1.15
–
0.88
–
0.60
–
0.53
0.43
–
–
–
–
–
0.50
0.75
2.00
3.00
4.00
5.00
5.50
6.00
7.00
8.00
9.00
10.00
11.00
12.00
–
–
–
0.90
–
–
0.26
0.34
0.86
1.19
1.67
2.04
–
2.38
2.86
–
3.57
3.66
–
–
–
–
–
0.29
0.39
0.86
1.27
1.70
2.09
–
2.40
2.60
–
3.58
3.84
–
–
–
1.20
3.00
4.16
5.83
7.14
–
8.33
10.00
–
12.50
12.80
–
–
–
1.72
1.90
2.08
2.16
2.27
–
2.64
2.70
2.70
2.67
2.67
2.75
2.69
2.67
5.0
5.0
5.1
4.5
4.3
–
3.9
3.0
2.7
2.5
2.3
2.2
2.0
1.3
6.60
6.00
5.49
4.39
4.40
–
3.33
2.49
2.20
2.27
2.00
1.67
1.50
1.10
2.19
3.13
3.85
4.23
4.33
4.64
5.00
5.00
5.36
5.36
5.45
5.45
–
0.62
0.89
1.10
1.20
1.24
1.33
1.43
1.43
1.53
1.53
1.56
1.56
–
0.72
0.98
1.11
1.12
1.19
1.22
1.43
1.31
1.40
1.54
1.56
1.64
–
9.0
10.0
11.0
12.0
12.5
15.0
20.0
–
–
–
–
–
–
–
[Ce(ClO4)4] = 4.0 × 10−4 M (for *), [HClO4] = 0.75 M, (A) [DMT] = 4.0 × 10−3 M (for #), [IrCl3] = 5.0 × 10−7 M (for and #). (B) [DET] = 1.0 × 10−3 M,
*
[IrCl3] = 2.0 × 10−7 M, (for * and #), for Bꢀ: kgr = 10−3 min−1
.
straight line passing through the origin tends to become par-
allel to the x-axis at higher concentrations (Fig. 2). Order of
the reaction shows first order kinetics at low concentrations
of organic substrates, which tends to become zeroth order
at their higher concentrations. This trend becomes clear
on plotting −dc/dt values versus [ketone] where straight
line passing through the origin tends to become parallel
to the x-axis at higher concentrations of organic substrates
(Fig. 3). In Table 1, it is seen that −dc/dt values increase
proportionately with increasing concentrations of organic
substrates in the beginning but at higher concentrations
the increase is not so prominent. kgr and kcalc values in
Table 1 for organic substrate variation are constant only at
low concentrations of organic substrates, while at higher
concentrations a increasing trend is obtained in both these
values. Proportionate increase in −dc/dt, kgr and kcalc values
for more than ten-fold variation in catalyst concentration
and fair constancy in kmolar values obtained for molar con-
centration of the catalyst {(kmolar = kgr/[IrCl3]) = 2.43 0.14
and 2.83 0.16 for dimethyl ketone and diethyl ketone,
respectively} indicate that the reaction follows first order
kinetics with respect to iridium(III) chloride concentrations
(Table 2). On plotting double logarithmic graphs between
log kgr and log[IrCl3] straight lines with slope values 0.926
and 1.06 were obtained (Fig. 4), which further confirm
Fig. 3. Effect of variation of [organic substrate] on the reaction rate at 30 ◦C.
(A and B) [Ce(ClO4)4] = 4.0 × 10−4 M, [HClO4] = 0.75 M, [IrCl3] = A – 5.0,
B – 2.0 (×10−7 M).
Fig. 2. Effect of variation of cerium(IV) concentrations on the rate at 30 ◦C.
A and B indicate same things as in Fig. 1 except for [Ce(ClO4)4].

86
P.K. Tandon et al. / Journal of Molecular Catalysis A: Chemical 232 (2005) 83–88
Table 2
Variation of [iridium(III)] and [H+] on the reaction rate at 25 ◦C
[IrCl3] × 107 M
kgr × 102 min−1
kcal × 102 min−1
[HclO4] M
kgr × 102 min−1
kcal × 102 min−1
A
B
A
B
A
B
A
B
1.0
2.0
3.0
4.0
5.0
6.0
7.0
8.0
9.0
–
–
0.32
0.59
0.82
1.13
1.43
1.62
1.96
–
2.37
2.77
3.08
–
–
0.34
0.58
0.86
1.07
1.42
1.59
1.99
–
2.45
2.84
3.05
–
0.50
0.60
0.75
1.00
1.50
2.00
2.50
3.00
3.50
4.00
4.50
5.00
1.56
1.43
1.10
0.98
0.91
1.05
1.16
1.42
1.71
1.84
0.71
0.58
0.55
0.59
0.66
0.82
1.53
1.49
1.09
0.92
0.96
0.93
1.18
1.08
1.28
1.88
0.69
0.59
0.57
0.69
0.64
0.73
1.18
0.40
0.58
0.72
1.11
1.49
–
2.02
2.31
2.50
2.88
3.18
0.71
0.91
1.10
1.43
–
2.04
2.32
2.50
2.86
3.07
1.43
1.59
1.79
2.00
0.99
1.11
1.14
1.14
10.0
11.0
12.0
(A) [Ce(ClO4)4] = 4.0 × 10−4 M, [DMT] = 4.0 × 10−3 M, [IrCl3] = 5.0 × 10−7 M, [HClO4] = 0.75 M. (B) [Ce(ClO4)4] = 4.0 × 10−4 M, [DET] = 1.0 × 10−3 M,
[IrCl3] = 2.0 × 10−7 M, [HClO4] = 0.75 M.
direct proportionality of the reaction velocity with respect to
[IrCl3]. Table 2 indicates a peculiar nature in case of perchlo-
ric acid variation where it is seen that −dc/dt, kgr and kcalc
values decrease in the beginning but after coming to a min-
imum the rate values start increasing with increasing [H+].
This nature becomes clear from Fig. 5 where a dip in the
graph is obtained in the beginning. Large volumes of sodium
perchlorate 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 ‘µ’
on the reaction rate was studied separately. Change in ionic
strength of the medium with the help of a standard solution
of sodium perchlorate does not affect the reaction velocity
and the rate values (−dc/dt) remain constant at 6.06 0.11
and 7.62 0.17 for dimethyl ketone and diethyl ketone, re-
spectively.
Fig. 5. Effect of variation of [H+] on the rate at 25 ◦C. (A)
[Ce(ClO4)4] = 4.0 × 10−4 M, [IrCl3] = 5.0 × 10−7 M, [DMT] = 4.0 × 10−3
M; (B) [Ce(ClO4)4] = 4.0 × 10−4 M, [IrCl3] = 2.0 × 10−7 M, [DET] = 1.0
× 10−3 M.
4. Discussion
Perchloric acid is a strong oxidant but below 50% concen-
tration and temperatures up to 50–60 ◦C, there is no release
of oxygen, thus, it will not act as oxidant under the conditions
in which the study was performed. Cerium(IV) in perchlo-
ric acid exists [8] in hydrolysed and unhydrolysed forms as
according to the following equilibrium:
K1
Ce4+ + H2O ꢀ Ce(OH)3+ + H+
K2
Ce(OH)3+ + H2O ꢀ Ce(OH)22+ + H+
Predominant species of cerium perchlorate in aqueous
perchloric acid medium is monomeric [9]. Increasing the
concentration of H+ ions from 0.1 to 4.0 M, concentration
of unhydrolysed species increases continuously while the
concentration of hydrolysed species first increases and then
Fig. 4. Effect of variation of [catalyst] on the rate at 30 ◦C.
[Ce(ClO4)4] = 4.0 × 10−4 M, [HClO4] = 0.75 M, A – [DMT] = 4.0, B –
[DET] = 1.0 (×10−3 M).

P.K. Tandon et al. / Journal of Molecular Catalysis A: Chemical 232 (2005) 83–88
87
the energy of activation values that the ease of formation of
the activated complex also lies in the above sequence. A sim-
ilar trend is observed with the entropy values also. Nearly
constant values of free energy of activation indicate that a
similar mechanism is operative in the oxidation of both the
cycloketones.
3−
It is known that IrCl3 in HCl gives IrCl6 species [11].
It has also been reported that iridium(III) and iridium(I) ions
are the stable species of iridium [12]. Furthermore, the aqua-
tion of [IrCl6]3− gives [IrCl5H2O]2−, [IrCl4(H2O)2]− and
[IrCl3(H2O)3] species [13–15] as shown by the following
equilibrium (2)
Scheme 1.
IrCl63− + n-H2O ꢀ [IrCl6−n(H2O)n]3−n + Cl−
(2)
from the vicinity of 1.0 M starts decreasing [10]. Our re-
sults, in the range of acid concentration where the studies
were performed, indicate the presence of both hydrolysed
as well as unhydrolysed species at low acid concentrations,
while at higher concentrations of perchloric acid more and
more hydrolysed species gets converted into the unhydrol-
ysed species. From Fig. 5 it is clear that dip in graphs corre-
sponds closely to the concentration from where concentration
of hydrolysed species starts decreasing. This is probably due
to the reason that hydrolysed species of cerium(IV) present
at low acid concentrations, being more reactive, reacts with
H+ ions giving rise to unhydrolysed species, which ultimately
takes part in the reaction as given in Scheme 1. Rate decreases
duetotheconversionofhydrolysedspeciesintotheunhydrol-
ysed species, which takes up hydrogen ions at low acid con-
centrations. With increasing [H+] this conversion slows down
and removal of H+ ions from the reaction mixture stops and
normal accelerating effect of the H+ ions starts taking place.
It was also observed, in ceric perchlorate and perchloric acid
variations, that initial titre values were exceptionally high at
low cerium(IV) and high acid concentrations, e.g. deviation
in the initial titre values from those of the calculated ones,
in case of dimethyl ketone and diethyl ketone, was found
to be 43.08–14.79% and 22.2–9.79%, respectively, from the
lowest to the highest cerium(IV) concentrations. All most
same increase ( 2–3%) in zero readings was obtained even
on performing the blank sets at similar concentrations with-
out adding the organic substrate and catalyst. Indicating the
probability of reduction of cerium(IV) to cerium(III) by wa-
ter, which is rapid at higher acid concentrations. Reduction
of cerium(IV) to cerium(III) may be checked if initial ceric
concentration is high but high ceric concentrations favour the
formation of unreactive polynulear complexes. Possibly high
acidity decreases the amount of hydrolysed species, which do
not react with water. Thus, the study was performed at 0.75 M
acid (except in acid variation) and maximum possible ceric
concentration, to minimize the reduction of cerium(IV) to
cerium(III) by water. The rate of reaction was found to be in
the order: diethyl ketone > dimethyl ketone. This sequence is
also confirmed from the energy of activation, entropy of acti-
vation and free energy of activation values, which were found
to be 12.76 and 9.15, −26.66 and −3993, 20.71 and 21.05
for dimethyl and diethyl ketones, respectively. It is clear from
In our study no effect of chloride ions on the reaction rate
was observed, which indicates that the above equilibrium is
shifted more towards the left side and IrCl5(H2O)2− can-
not be considered as the reactive species [4,16]. Therefore,
3−
considering our experimental results, IrCl6 has been con-
sidered to be the reactive species of iridium(III) chloride in
the present study. Thus, according to the proposed mech-
anism (Scheme 1), cerium(IV) species combines with or-
ganic 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 takes up a hydro-
gen ion giving rise to iridium(I) species and the intermedi-
ate product. Iridium(I) species in the fast step takes up two
cerium(IV) molecules to regenerate the original iridium(III)
species while the intermediate product is further oxidized in
the fast step to the final oxidation products.
Formation of a 1:1 complex between cerium(IV) and alco-
hols [17] and ketones [18] is well documented. Complex for-
mation between cerium(IV) and alcohols to give complexes
of composition [ROH·cerium(IV)]4+ and Michaelis–Menten
type of kinetics has been reported [19]. While similar re-
sults have been reported in the case of ketones also [20,21].
Formation of a complex between cerium(IV) and the organic
substrate as obtained from our data is supported by the change
from first order to zero order kinetics and also from the con-
stancy in the −dc/dt, kgr and kcalc values only in the beginning
when the complex formation is small. Deviations from con-
stancy become more pronounced at higher [CeIV] or [sub-
strate]. Considering equilibrium concentrations in steps (I)
and (II) of the mechanism and putting concentration of C1
from step (I) into the concentration of [IrIII] obtained from
step (II), total concentration of catalyst may be given as ac-
cording to Eq. (3)
[C2]
K1K2[S][CeIV]
[IrIII]total
=
+ [C2]
(3)
From Eq. (3) concentration of complex C2 is given as
K1K2[CeIV][S][IrIII]T
[C2] =
(4)
1 + K1K2[CeIV]

88
P.K. Tandon et al. / Journal of Molecular Catalysis A: Chemical 232 (2005) 83–88
Rate in terms of decreasing concentration of cerium(IV) from
step (III) of the mechanism may be given as
readings at low oxidant or high acid concentrations and
reduction of cerium(IV) even in the absence of organic
substrate indicate for this possibility, but instability of
aqueous Ceric perchlorate solution in aqueous perchloric
acid solution as reported by many workers may also result
in the high initial readings. However, even the existence of
this possibility will not affect the final rate law. Thus we can
safely assume the validity of the final rate law and at least
formation of the complexes before the rate determining step.
d[CeIV]
2kK1K2[CeIV][S][IrIII][H+]
1 + K1K2[CeIV][S]
−
=
(5)
dt
Equation is multiplied by 2 because two moles of cerium(IV)
are required to regenerate the catalyst in its original form.
This equation explains all experimental findings except the
nature shown by the [H+] ions at their low concentrations.
Probable reason for the initial decrease in rate values at low
acid concentrations has already been discussed in the begin-
ning. At low concentrations of oxidant and organic substrate
the inequality 1 ꢀ kK1K2[CeIV][S] may hold and Eq. (5) re-
duces to Eq. (6) which, explains the nature shown by various
reactants.
Acknowledgement
S.S. and A.K.S. (C.S.T., U.P. D-3205) and M.P. (U.G.C.
F.12-97/2001,SR-I) are grateful for the financial assistance.
d[CeIV]
−
= 2kK1K2[CeIV][S][IrIII][H+]
(6)
dt
References
At higher concentrations of oxidant and substrate the reverse
inequality 1 ꢁ kK1K2[CeIV][S] holds good and the equation
becomes
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−d[CeIV]/dt
2kK1K2[CeIV][S]
1 + K1K2[CeIV][S]
= kꢂ =
(7)
[IrIII]T[H+]
where kꢂ is another constant. Further verification of the rate
law (5) may be given by rewriting Eq. (8) as
1
1
kꢂ
1
1
=
= +
2kK1K2[CeIV][S] 2k
(8)
Vi
From this equation value of k from the intercepts between 1/kꢂ
versus 1/[substrate] and 1/kꢂ versus 1/[cerium(IV)], comes
out to be 16.0; 21.05 and 62.50; 62.50, while the value of
kK1K2 from the slopes comes out to be 5.0; 4.17 (×106) and
1.80; 1.60 (×106) for dimethyl ketone and diethyl ketone,
respectively. Fair constancy in the rate values calculated from
two graphs further indicates the validity of Scheme 1 and the
rate law (5).
Possibility of interaction between the organic substrate
and iridium giving up the complex C1 in first step of
Scheme 1, can be ruled out as no reaction between these two
in the absence of oxidant was observed. No effect of chloride
ions on the rate negates the possibility of their release before
the rate determining step. Thus, the only alternative may be
that cerium(IV) reacts with iridium(III) in the first step to
form a complex, which in turn reacts with organic substrate
to give the complex C2. Exceptionally high values of initial
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