Kinetics and mechanism of the Ir(III)-catalyzed oxidation of xylose and maltose by potassium iodate in aqueous alkaline medium

Article abstract of DOI:10.1016/j.carres.2007.02.010

For the first time, the Ir(III) catalysis of the iodate oxidation of xylose and maltose in aqueous alkaline medium has been investigated. The reactions exhibit first-order kinetics with respect to lower [IO₃^-] and [OH^-] and show zero-order kinetics at their higher concentrations. Unity order at low concentrations of maltose becomes zero order at its higher concentrations, whereas zero-order kinetics with respect to [xylose] was observed throughout its variation. The reaction rate is found to be directly proportional to [Ir(III)] in the oxidation of both reducing sugars. Negligible effect of [Cl^-] and nil effect of ionic strength (μ) on the rate of oxidation have also been noted. The species, [IrCl₃(H₂O)₂OH]^- was ascertained as the reactive species of Ir(III) chloride for both the redox systems. Various activation parameters have been calculated. Formic acid and arabinonic acid for maltose and formic acid and threonic acid for xylose were identified as the main oxidation products of the reactions. Mechanisms consistent with the observed kinetic data and spectral evidence have been proposed for the oxidation of xylose and maltose.

Full text of DOI:10.1016/j.carres.2007.02.010

Carbohydrate Research 342 (2007) 1078–1090
Kinetics and mechanism of the Ir(III)-catalyzed oxidation of xylose
and maltose by potassium iodate in aqueous alkaline medium
*
Ashok Kumar Singh, Shalini Srivastava, Jaya Srivastava and Reena Singh
Department of Chemistry, University of Allahabad, Allahabad 211 002, India
Received 29 December 2006; received in revised form 1 February 2007; accepted 8 February 2007
Available online 16 February 2007
Abstract—For the first time, the Ir(III) catalysis of the iodate oxidation of xylose and maltose in aqueous alkaline medium has been
ꢀ
ꢀ
investigated. The reactions exhibit first-order kinetics with respect to lower ½IO ꢁ and [OH ] and show zero-order kinetics at their
3
higher concentrations. Unity order at low concentrations of maltose becomes zero order at its higher concentrations, whereas zero-
order kinetics with respect to [xylose] was observed throughout its variation. The reaction rate is found to be directly proportional to
ꢀ
[
Ir(III)] in the oxidation of both reducing sugars. Negligible effect of [Cl ] and nil effect of ionic strength (l) on the rate of oxidation
ꢀ
have also been noted. The species, [IrCl
3
(H
2
O)
2
OH] was ascertained as the reactive species of Ir(III) chloride for both the redox
systems. Various activation parameters have been calculated. Formic acid and arabinonic acid for maltose and formic acid and thre-
onic acid for xylose were identified as the main oxidation products of the reactions. Mechanisms consistent with the observed kinetic
data and spectral evidence have been proposed for the oxidation of xylose and maltose.
ꢀ
2007 Elsevier Ltd. All rights reserved.
Keywords: Mechanism; Xylose; Maltose; Potassium iodate; Alkaline medium; Ir(III) Catalysis
1
5
1
. Introduction
odic acid have been investigated in acidic media. Re-
ports regarding the use of periodate in the uncatalyzed
oxidation of carbohydrates and polymeric substrate
and Ru(III)- and ruthenate ion-catalyzed oxidation of
reducing sugars in alkaline medium are also avail-
The study of carbohydrates is one of the most exciting
fields of organic chemistry. Vast literature is available
on the kinetics of oxidation of carbohydrates by various
organic and inorganic oxidants. The oxidation of
aldoses by chlorine, bromine and iodine have been
reported in alkaline media. The aldonic acids as
primary products of oxidation of aldoses by bromine
have been extensively studied by Isbell and co-workers,
who pointed out that b-aldoses (C-1 equatorial) are oxi-
dized much faster than a-aldoses (C-1 axial). The cata-
lyzed and uncatalyzed oxidation of sugars have been
studied in detail by using N-halo compounds.
ganic oxidants such as Cu(II), ammonical Ag(I) and
Nessler’s reagent have been used in the uncatalyzed oxi-
dation of sugars in aqueous alkaline medium.
mechanisms for the oxidations of some aldoses by
Cr(VI), V(V), Ce(IV), Mn(III), Ir(IV), Au(III) and peri-
1
6–18
able.
The uncatalyzed oxidation of oxalic acid, ace-
tophenone and benzaldehydes and Ru(III)-catalyzed
oxidation of a-hydroxy acids by iodate in acidic medium
In the reported Ru(III)- and
Os(VIII)-catalyzed oxidation of styrene and stillbene in
ꢀ
acidic medium, IO3 has also been used as an oxidant.
But as of this date no report is available, in which iodate
as an oxidant and Ir(III) chloride as homogeneous cata-
lyst have been used in the oxidation of monosaccharides
and disaccharides in alkaline medium. In view of the
biological importance of reducing sugars, and also in
view of the fact that no investigation on the kinetic oxi-
dation of reducing sugars with iodate as an oxidant and
Ir(III) chloride as homogeneous catalyst in alkaline
medium has so far been made, the present study has
been undertaken. In this study our main aim was to
ascertain the following:
1
1
9–22
have been reported.
2
23
3
–11
Inor-
1
2–14
The
*
Corresponding author. Tel.: +91 0532 2640434; e-mail: ashokeks@
rediffmail.com
0
008-6215/$ - see front matter ꢀ 2007 Elsevier Ltd. All rights reserved.
doi:10.1016/j.carres.2007.02.010

A. K. Singh et al. / Carbohydrate Research 342 (2007) 1078–1090
1079
ꢀ
(
(
1) IO3 as an oxidant in alkaline medium behaves in
solution to the reaction mixture, and the progress of the
reaction was monitored by iodometrically estimating the
amount of unreacted iodate at regular time intervals.
The unreacted iodate in presence of excess potassium io-
dide and perchloric acid is converted into iodine, and the
liberated iodine is titrated against a standard solution of
sodium thiosulfate using starch solution as an indicator.
the same way as it behaved in the reported
Ru(III)-catalyzed oxidation of indigo carmine by
iodate in aqueous sulfuric acid solution.
2
4
2) The role of reducing sugar molecule in the presence
ꢀ
of Ir(III) chloride as catalyst and IO3 as an oxi-
dant in alkaline medium is similar to the reported
role of reducing sugar in the presence of Ir(III)
chloride and N-bromoacetamide (NBA) in acidic
2.2. Stoichiometry and product analysis
6
medium.
ꢀ
(
(
3) Ir(III) Chloride in alkaline medium participates in
the reaction under investigation in the same way,
as it is reported for Ir(III)-catalyzed oxidation of
A set of solutions of varying ½IO ꢁ:[reducing sugar]
3
ratios were equilibrated at room temperature for 72 h
ꢀ
under the conditions of ½IO ꢁ ꢂ ½sugarꢁ. Estimation
3
6
ꢀ
reducing sugars by NBA in acidic medium.
of residual ½IO ꢁ in different sets showed that 1 mole
3
ꢀ
4) There is any possibility of the formation of a com-
plex between the reactive species of Ir(III) chloride
and the reactive species of potassium iodate in
of sugar consumed 2 and 4 mol of IO for xylose and
3
maltose, respectively. Accordingly, the following stoichio-
metric equations were formulated:
–
Ir(III)/OH
–
–
C H O + 2IO3
C H O
5
+ HCOOH + 2IO2
5
10
5
4
8
D-xylose
Threonic acid Formic acid
–
Ir(III)/OH
–
–
C H O + 4IO3
2C H O + 2HCOOH + 4IO2
1
2
22 11
5
10
6
Maltose
Arabinonic acid Formic acid
alkaline medium in the same form as it is reported
in the Ir(III) chloride-catalyzed oxidation of reduc-
ing sugars by NBA in acidic medium.
Formic acid and arabinonic acid in the case of maltose
and formic acid and threonic acid in the case of xylose
were identified as the main oxidation products by the
help of equivalence and kinetic studies and also by
TLC and conventional spot test methods.
6
(
5) The formation of an activated complex results by
the interaction of two oppositely charged species
in the same way as it is reported for Ir(III) catalyzed
oxidation of reducing sugars by NBA in acidic
6
medium.
3. Results
The oxidation of xylose and maltose were investigated at
different initial concentrations of the reactants. The time
course of the reaction in the oxidation of both reducing
sugars is shown in Figure 1. In the case of xylose, the ini-
tial rate of the reaction, that is, (ꢀdc/dt) in each kinetic
run, was calculated from the slope of the tangent of the
2
. Experimental
2
.1. General procedures
A solution of Ir(III) chloride (Johnson–Mathey & Co.
Ltd) was prepared in HCl of known strength
ꢀ
ꢀ
plot of unconsumed ½IO ꢁ versus time at fixed ½IO3 ꢁ,
3
ꢀ2
ꢀ
(
10.00 · 10 mol/L) and its concentration was main-
except in ½IO ꢁ variation when a tangent has been
3
ꢀ
3
tained at 6.69 · 10 mol/L. A standard stock solution
drawn at a fixed time. In the case of maltose, the initial
rate of reaction was calculated by the slope of the
straight line obtained by the plot of unconsumed
of KIO (E. Merck) was prepared by dissolving a known
3
weight of KIO in double-distilled water, and its concen-
3
ꢀ
tration was ascertained iodometrically. Standard solu-
tions of xylose and maltose (A.R. grade) were freshly
prepared. NaOH and KCl (E. Merck) were employed
to maintain the required alkalinity and ionic strength,
respectively. All the reactions were studied at constant
temperature 40 ꢁC (± 0.1 ꢁC). The reaction was initiated
by adding the requisite volume of pre-equilibrated sugar
½IO ꢁ versus time (Fig. 1). For the determination of
3
the order of reaction with respect to each reactant, the
help of Ostwald’s isolation method in conjunction with
van’t Hoff’s differential method has been taken. Consid-
ꢀ
ering IO3 , reducing sugar, hydroxyl ions and Ir(III)
chloride as the main reactants, the general form of rate
equation for the reaction can be written as

1080
A. K. Singh et al. / Carbohydrate Research 342 (2007) 1078–1090
ꢀ
Table 1. Effect of variations in ½IO ꢁ and ionic strength (l) of the
2
2
1
1
5
0
5
0
5
0
3
medium on the rate of oxidation of xylose and maltose at 40 ꢁC
4
ꢀ1
½
IO3ꢀꢁ ꢃ 103 (mol/L)
l (mol/L)
k
1
· 10 (s
)
Xylose
Maltose
0
0
1
1
1
2
2
2
3
3
4.00
1
1
1.00
1.00
1
.40
.80
.00
.20
.60
.00
.40
.80
.20
.60
0.25
0.25
0.25
0.25
0.25
0.25
0.25
0.25
0.25
0.25
0.25
0.25
0.50
0.75
1.00
1.25
5.95
6.25
6.00
6.67
6.25
6.36
6.25
5.95
—
4.44
4.44
5.00
4.75
5.00
5.00
4.69
2.08
1.95
2.00
1.98
1.93
1.97
1.89
1.51
1.32
—
1.04
2.22
2.24
2.22
2.24
2.16
(
1)
(
2)
0
20
40
60
80
100
.00
.00
Time (min)
ꢀ
Figure 1. Plots between remaining ½IO ꢁ and time at 40 ꢁC: (1)
3
ꢀ
ꢀ3
ꢀ
ꢀ2
½
IO ꢁ ¼ 2:00 ꢃ 10 mol=L, [OH ] = 20.00 · 10 mol/L, [Ir(III)] =
3
ꢀ
6
ꢀ2
.00
1
3.38 · 10 mol/L, [maltose] = 2.00 · 10 mol/L, [l] = 0.25 mol/L.
ꢀ
ꢀ3
ꢀ
ꢀ2
(
2
2) ½IO ꢁ ¼ 1:20 ꢃ 10 mol=L, [OH ] = 20.00 · 10 mol/L, [Ir(III)] =
ꢀ
ꢀ2
3
Solution conditions: [OH ] = 20.00 · 10 mol/L, [Sugar] = 2.00 ·
ꢀ
6
ꢀ2
6.76 · 10 mol/L, [xylose] = 2.00 · 10 mol/L, [l] = 0.25 mol/L.
ꢀ2
ꢀ6
1
1
0
0
mol/L, [Ir(III)] = 26.76 · 10 mol/L (xylose) and 13.38 ·
ꢀ6
mol/L (maltose).
rate ¼ k½IO ꢀꢁ ½sugarꢁ ½IrðIIIÞꢁ ½OH ꢁ
a
b
c
ꢀ d
ð1Þ
3
1
1
1
8
5
2
9
6
3
0
For the determination of the experimental rate law, first
of all a series of experiments with varying initial concen-
tration of IO3 were performed at constant concentra-
xylose
ꢀ
tions of all other reactants and at constant
temperature 40 ꢁC. In the light of Ostwald’s isolation
method, the concentrations of reducing sugar and
ꢀ
ꢀ
OH ions were fixed in large excess with respect to
IO3 throughout its variation. Since the concentrations
of those in excess will not change very much during
the course of reaction and Ir(III) being a catalyst is
reproduced in the reaction, the rate law (Eq. 1) under
this condition becomes
maltose
rate ¼ k ½IO3ꢀꢁ
a
ð2Þ
=
1
where k , the apparent rate constant,
k[sugar] [Ir(III)] [OH ] .
Kinetic results obtained with the variation of IO
to 4.0 · 10 mol/L are presented in
Table 1 and also in Figure 2. It is clear from Figure 2
1
b
c
ꢀ d
0
1
2
3
4
5
ꢀ
3
[IO3 ]×10 (mol L–1)
–
3
ꢀ
3
ꢀ3
from 0.4 · 10
ꢀ
ꢀ
Figure 2. Plots between (ꢀdc/dt) and ½IO ꢁ at 40 ꢁC: [OH ] =
3
ꢀ
ꢀ
2
ꢀ6
that in the oxidation of both xylose and maltose, there
20.00 · 10 mol/L, [Ir(III)] = 26.76 · 10 mol/L (xylose) and
6
ꢀ2
ꢀ
13.38 · 10 mol/L (maltose), [Sugar] = 2.00 · 10 mol/L, [l] =
.25 mol/L.
is direct proportionality between the rate and ½IO
ꢁ
3
0
throughout its variation except at higher concentrations,
where the rate is unaffected by the ½IO3ꢀꢁ. This shows
ꢀ
that the order of reaction with respect to ½IO ꢁ is unity
3
ꢀ3
ꢀ
up to its 2.40 · 10 mol/L, and thereafter it becomes
with respect to ½IO ꢁ has been taken as unity. The pseudo
3
zero order. Almost constant values of the pseudo first-
first-order rate constant, k , was calculated as
1
ꢀ3
order rate constant (k ) up to its 2.40 · 10 mol/L also
support the above experimental findings, where unity
order with respect to IO3 has been concluded (Table
1
ðꢀdc=dtÞ
k
¼
ð3Þ
1
ꢀ
½
IO
ꢁ
ꢀ
3
1
). Since throughout the study of variations of all other
After establishing the first-order kinetics with respect to
ꢀ
ꢀ
reactants concentrations, the concentration of IO3 was
fixed in its lower range; hence, for the purpose of calcu-
lation of pseudo-first-order rate constant, k , the order
½IO ꢁ in its lower range, nearly ten-fold variations in the
3
concentrations of xylose and maltose were made under
pseudo first-order conditions and on the basis of
1

A. K. Singh et al. / Carbohydrate Research 342 (2007) 1078–1090
1081
observed kinetic data, zero-order kinetics in [xylose]
8
6
4
2
0
throughout its variation and first-order kinetics in [malt-
ose] in its lower range were observed (Figs. 3 and 4). It is
also observed that the reaction rate becomes indepen-
dent of maltose at its higher concentrations (Fig. 4).
Since in an unusual way, unity order in lower maltose
concentration was observed; hence, to verify the kinetic
results, experiments for the variations of [maltose] have
been performed under three different conditions at
(3)
(
2)
4
4
0 ꢁC. The results thus obtained are presented in Figure
. It is apparent from Figure 4 that there is first-order
(
1)
kinetics at low [maltose], which becomes zero order at
higher [maltose]. Unity order in [Ir(III)] up to nearly
1
0-fold variation was observed in the oxidation of both
the reducing sugars (Fig. 5). At low concentrations the
pseudo first-order rate constant (k ) is found to be di-
0
2
4
6
8
10
12
1
ꢀ
rectly proportional to [OH ], and at higher concentra-
2
–1
[
Maltose]×10 (mol L )
ꢀ
tions it becomes almost independent of [OH ] in the
ꢀ
Figure 4. Plots between k
1
and [maltose] at 40 ꢁC: (1) ½IO ꢁ ¼
oxidation of both xylose and maltose (Fig. 6). When
variation in ionic strength of the medium was made, it
was observed that there is almost no change in pseudo-
3
ꢀ
3
ꢀ
ꢀ2
0
1
1
:80 ꢃ 10 mol=L, [OH ] = 15.00 · 10 mol/L, [Ir(III)] = 20.07 ·
ꢀ
ꢀ2
6
ꢀ
ꢀ3
ꢀ
0
0
mol/L. (2) ½IO ꢁ ¼ 1:00 ꢃ 10 mol=L, [OH ] = 10.00 ·
3
ꢀ6
ꢀ
mol/L, [Ir(III)] = 13.38 · 10 mol/L. (3) ½IO ꢁ ¼ 1:00 ꢃ
3
ꢀ3
ꢀ
ꢀ2
ꢀ6
first-order rate constant, k , values with the change in
10 mol=L, [OH ] = 20.00 · 10 mol/L, [Ir(III)] = 13.38 · 10 mol/L.
1
ionic strength of the medium (Table 1). Since there
ꢀ
was no effect of [Cl ] on the rate of oxidation, the ionic
strength of the medium was fixed by the addition of KCl
1
0
8
6
4
2
0
to the reaction mixture. Observed k values at 30, 35, 40
1
xylose
#
#
#
and 45 ꢁC were utilized to calculate E , DH , DS , DG
a
maltose
and the Arrhenius frequency factor (A) in the oxidation
of aforesaid reducing sugars (Table 2). On the basis of
observed first-order kinetics in lower concentrations of
ꢀ
ꢀ
IO3 and OH , zero order in [xylose], first order in low-
er concentrations of maltose and first order in [Ir(III)],
the following experimental rate laws in the form of
8
6
4
2
0
0
10
20
30
40
50
60
Ir(III)]×10 (mol L–1)
6
[
ꢀ
Figure 5. Plots between
k
1
and [IrCl
3
]
at 40 ꢁC: ½IO3 ꢁ ¼
ꢀ3
ꢀ
ꢀ2
1
1
:00 ꢃ 10 mol=L, [OH ] = 20.00 · 10 mol/L, [Sugar] = 2.00 ·
ꢀ
2
0
mol/L.
Eqs. 4 and 5 can be proposed for the oxidation of xylose
and maltose, respectively.
d½IO ꢀꢁ
3
¼ k ½OH ꢁ½IO ꢀꢁ½IrðIIIÞꢁ
0
ꢀ
ꢀ
ð4Þ
3
dt
and
d½IO ꢀꢁ
0
2
4
6
8
10
12
3
¼ k ½OH ꢁ½IO ꢀꢁ½IrðIIIÞꢁ½Sꢁ
00
ꢀ
ꢀ
ð5Þ
2
–1
3
dt
[
xylose]×10 (mol L
)
0
00
ꢀ
where ‘S’ stands for maltose in Eq. 5 and k and k are
the composite rate constants for xylose and maltose
respectively.
Figure 3. Plot between k
1
and [xylose] at 40 ꢁC: ½IO ꢁ ¼ 1:00ꢃ
3
ꢀ3
ꢀ
ꢀ2
ꢀ6
10
mol=L, [OH ] = 20.00 · 10 mol/L, [Ir(III)] = 26.76 · 10
mol/L.

1082
A. K. Singh et al. / Carbohydrate Research 342 (2007) 1078–1090
2
8
6
1
1
2
0
8
6
4
2
0
of N-bromosuccinimide (NBS) and NBA in presence
of Ir(III) as the homogeneous catalyst. In each case, the
existence of the following equilibrium:
maltose
xylose
3ꢀ
2ꢀ ꢀ
5 2
(H O)] + Cl
[
IrCl
6
]
+ H
2
O ꢀ [IrCl
ꢀ
has been assumed, and on the basis of the effect of [Cl ]
on the rate of reaction, IrCl6
3ꢀ
is taken as the reactive
species of Ir(III) chloride in acidic medium. Very
recently a study concerning the separation of rhodium
and iridium and chloridation and chlorination of Ir(III)
2
9
and Ir(IV) is reported. In this paper authors have
made efforts to study the effect of HCl concentration
and temperature on Ir(III) speciation at equilibrium.
According to them, at room temperature and at 70 ꢁC,
Ir(III) chloride in 0.1 M HCl concentration will remain
as [IrCl (H O) ], whereas in 8 M HCl solution, it will
0
20
40
60
–
2
–1
[
OH ]×10 (mol L )
3
2
3
2
ꢀ
3ꢀ
remain as [IrCl (H O)]
and [IrCl6] , respectively.
5
2
ꢀ
ꢀ
Figure 6. Plots between
k
1
and [OH ] at 40 ꢁC: ½IO ꢁ ¼
3
ꢀ3
ꢀ2
Between concentrations of 0.1 and 8 M HCl,
[IrCl (H O) ], [IrCl (H O) ] and [IrCl (H O)] , are
3 2 3 4 2 2 5 2
1
1
:00 ꢃ 10 mol=L, [Sugar] = 2.00 · 10 mol/L, [Ir(III)] = 26.76 ·
ꢀ
2ꢀ
ꢀ
6
ꢀ6
0
mol/L (xylose) and 13.38 · 10 mol/L (maltose), [l] = 0.55 M.
reported to be the predominant species. In the present
investigation, since the solution of the catalyst, Ir(III)
chloride, has been prepared in 0.1 M HCl, and in view
of the observed kinetic data and the reported literature,
it is reasonable to assume that the starting species of
Ir(III) chloride is [IrCl (H O) ]. Furthermore, since the
Table 2. The values of composite rate constant and activation
parameters for the oxidation of xylose and maltose at 40 ꢁC
Activation parameters
Xylose
Maltose
3
2
3
ꢀ
2
1
E
a
0
0
(kJ mol
)
61.57
58.77
—
study for the catalyzed oxidation of reducing sugars
has been made in alkaline medium, a decision about
the reactive species of Ir(III) chloride can be made only
ꢀ2
ꢀ1
ꢀ1
ꢀ1
ꢀ1
2
k (mol L s
)
)
1.12 · 10
0
ꢀ3
3
4
k
(mol L s
—
0.39 · 10
22.26
56.18
46.69
—
#
ꢀ1
DS (J K mol
)
)
)
ꢀ17.37
58.98
53.55
ꢀ
#
ꢀ1
after taking into account the effect of [OH ] on the rate
DH (J K mol
#
ꢀ1
ꢀ1
DG (J K mol
of oxidation. On the basis of (1) the observed kinetic
data relating to first-order kinetics at low concentrations
ꢀ2
2
ꢀ1
ꢀ1
12
A (mol L s
)
)
1.87 · 10
ꢀ3
3
14
ꢀ
A (mol L s
—
2.22 · 10
of OH in the oxidation of both xylose and maltose and
(2) an increase in absorbance from 1.68 to 2.08 and 2.44
of Ir(III) chloride solution and Ir(III) chloride solution
with two different concentrations of OH (Fig. 7: (2),
4. Discussion
ꢀ
(
3) and (4)), it can be concluded that the following equi-
4
.1. Reactive species of Ir(III) chloride in alkaline medium
librium is established:
ꢀ
ꢀ
2 2 2
O) ] + H O
A spectrophotometric study of the kinetics of the hydra-
[IrCl
3
(H
2
O)
3
] + OH ꢀ [IrCl
3
ꢄOH(H
3
ꢀ
ꢀ
tion of IrCl6
and of the addition of a Cl to
2
ꢀ
[
Ir(OH )Cl ] in 1.0–2.5 M HClO (or HCl) at 50 ꢁC
Out of two species, [IrCl
3
(H O)
2
3
]
and [IrCl
3
Æ
2
5
4
2
5
ꢀ
ꢀ
is reported, where evaluation of the rate constants
and the equilibrium constant K for the reversible
reaction
OH(H
2
O)
2
] , the species [IrCl
assumed as the reactive species of Ir(III) chloride in
the oxidation of xylose and maltose, because with the
reactive species [IrCl ÆOH(H O) ] , and taking into con-
sideration the existence of above equilibrium in the reac-
tion, a rate law can be derived which will show that there
3
ÆOH(H
2
O)
2
]
can be
ꢀ
3
2
2
k1
3
ꢀ
ꢁ2ꢀ þ Clꢀ
IrCl6 þ H
2
O ꢀ ½IrðH
2
OÞCl
5
kꢀ1
ꢀ
is a positive effect of [OH ] on the rate of oxidation.
have been made. Visible and ultraviolet absorption spec-
ꢀ
tra of the new Ir(III) complexes IrðH
2
OÞ Cl
and
and
2
4
3ꢀ
Ir(H O) Cl , together with the spectra of IrCl
4.2. Reactive species of potassium iodate in alkaline
medium
2
3
3
6
2ꢀ
IrðOH
2
ÞCl
in 2.5 F HClO –1.2 F NaClO are also
reported and found in reasonable agreement with the
results reported by Poulsen and Garner in water and
by Jørgensen. A few recent reports are available in
the literature, in which kinetic studies have been made
for the oxidation of reducing sugars by acidic solution
5
4
4
2
6
2
5
It is reported that KIO has been used as an oxidant in
3
2
7
20
30
the oxidation of acetophenones, ferrocyanide, thio-
cynate and 1,3-dihydroxybenzene in acidic medium.
In each case IO3 has been regarded as the reactive
3
1
32
ꢀ

A. K. Singh et al. / Carbohydrate Research 342 (2007) 1078–1090
1083
ꢀ
ꢀ
4
.5
3
IO3 , which very well supports the existence of
IO3 in alkaline medium. In view of the reported kinetic
data and spectral evidence, it can very easily be
concluded that the species IO3 is the reactive species
of KIO in the oxidation of xylose and maltose in alka-
line medium.
Series1
Series2
Series3
Series4
Series5
Series6
Series7
Series8
ꢀ
3
2
1
0
3
4.3. Reactive form of sugar in alkaline medium
3
3
It is reported that in the presence of alkali reducing
sugars undergo a tautomeric change resulting in the for-
mation of an enediol anion and an enediol. The base-
catalyzed formation of the enediol can be shown as
follows:
.5
2
(
a) Aldehydo sugars
.5
1
-
H
C
C
R
O
H
C
C
R
O
-
H
OH + OH
OH + H O
2
enediol anion
-
H
C
C
R
O
H
C
C
R
OH
OH + OH
.5
0
-
OH + H O
2
enediol anion
(1,2-endiol)
200
250
300
350
400
Wavelength (nm)
(
b) Keto sugars
ꢀ
ꢀ3
ꢀ5
Figure 7. Spectra: (1) ½IO ꢁ ¼ 1:00 ꢃ 10 mol=L. (2) [Ir(III)] =
3
ꢀ5
ꢀ
2
1
1
1
.68 · 10 mol/L. (3) [Ir(III)] = 2.68 · 10 mol/L, [OH ] = 20.00 ·
ꢀ2
ꢀ2
ꢀ2
ꢀ5
ꢀ5
ꢀ
ꢀ
0
0
0
mol/L. (4) [Ir(III)] = 2.68 · 10 mol/L, [OH ] = 50.00 ·
mol/L. (5) [Ir(III)] = 2.68 · 10 mol/L, [OH ] = 20.00 ·
H
H
C
C
R
OH
H
C
C
R
OH
O
ꢀ
ꢀ3
ꢀ5
mol/L, ½IO ꢁ ¼ 1:00 ꢃ 10 mol=L. (6) [Ir(III)] = 2.68 · 10
-
3
-
+
ꢀ
ꢀ2
ꢀ
ꢀ3
O
H O
2
mol/L, [OH ] = 20.00 · 10 mol/L, ½IO ꢁ ¼ 2:00 ꢃ 10 mol=L. (7)
+ OH
3
ꢀ
5
ꢀ
ꢀ2
ꢀ
[
Ir(III)] = 2.68 · 10 mol/L, [OH ] = 20.00 · 10 mol/L, ½IO ꢁ ¼
3
ꢀ3
ꢀ2
1
2
1
:00 ꢃ 10 mol=L, [maltose] = 2.00 · 10 mol/L. (8) [Ir(III)] =
enediol anion
ꢀ5 ꢀ ꢀ2
ꢀ
.68 · 10 mol/L, [OH ] = 20.00 · 10 mol/L, ½IO3 ꢁ ¼ 1:00 ꢃ
ꢀ3
ꢀ2
0
mol=L, [maltose] = 10.00 · 10 mol/L.
H
C
OH
H
C
C
R
OH
-
-
species of KIO in acidic medium. Kinetic studies for
3
C
R
O
+ H O
OH + OH
2
3
23,24
2
Os(VIII)- and Ru(III)-catalyzed
oxidation of or-
ganic compounds by an acidic solution of iodate are also
ꢀ
enediol anion
(
1,2-enediol)
reported. In these cases either HIO or IO3 has also
3
been concluded to be the reactive species of KIO in
3
acidic medium. Until now, no report is available in the
literature where KIO has been used as an oxidant in
3
either the catalyzed or uncatalyzed oxidation of organic
compounds in alkaline medium. However, a few reports
The formation of the enediol anion and the enediol in
the presence of alkali is also supported by the work of
3
4
are available where NaIO has been used as an oxidant
Isbell and co-workers.
In the present study the
4
1
7
18
ꢀ
in Ru(III)- and ruthenate ion-catalyzed oxidation of
reducing sugars in alkaline medium. On the basis of
equivalence and kinetic studies for both the redox pro-
observed order with respect to [OH ] for the Ir(III)-
catalyzed oxidation of reducing sugars has led us to
assume that it is the enediol form of sugar which is actu-
ally taking part in the reactions under investigation.
ꢀ
cesses, IO4 is reported to be finally converted into

1084
A. K. Singh et al. / Carbohydrate Research 342 (2007) 1078–1090
ꢀ
4.4. Spectral evidence for the formation of complexes
during the course of reaction
auxochrome, OH , to give rise to another chromophore,
2
–
O
O
.
I — O–IrCl (H O)OH
3
2
After making a decision about the reactive species of
both Ir(III) chloride and potassium iodate, and also
about the reactive form of reducing sugars, efforts have
been made to ascertain the possibility of the formation
of a complex or complexes during the course of reaction.
Contrary to the observed zero effect of [xylose] on the
rate of oxidation, first-order kinetics at low [maltose]
and zero-order kinetics at high [maltose] have been
observed in Ir(III)-catalyzed oxidation of maltose by
an alkaline solution of potassium iodate. By considering
the first-order kinetics in respect to the maltose concen-
tration, an effort was made to probe the possibilities of
the formation of a complex or complexes between
2
It is reported that transition metal ions such as Ru(III)
3
and Pd(II) form complexes with the reducing sugar
3
,4
molecule in acidic medium. It is also reported that
the complexes thus formed have tendencies to react with
reactive species of oxidant NBA and also with Hg(II),
ꢀ
ꢀ
IrCl (H O) OH] and IO3 and between the complex,
3 2 2
[
ꢀ
if any, formed due to interaction of reactive species of
the catalyst and reactive species of oxidant, and the reac-
tive form of reducing sugar molecule in the catalyzed
oxidation of maltose. Conclusions for the formation of
complexes in the reaction under investigation have been
whose function in the reaction was as a Br ion scaven-
ger and as co-catalyst. In the oxidation of xylose by an
alkaline solution of potassium iodate in presence of Ir-
(
III) as homogeneous catalyst, it has been observed that
order with respect to xylose concentration is zero
throughout its 10-fold variation. This shows that the
enediol form of xylose will not take part in the reaction
before the rate-determining step, although in fast steps,
it will combine with the most reactive complex to form
the reaction products along with regeneration of the cat-
alyst. It is thus clear that the possibility of the formation
of a complex in oxidation of xylose occurs only between
reactive species of Ir(III) chloride and reactive species
of potassium iodate in alkaline medium. In order to
probe the possible formation of a complex between
ꢀ
drawn on the basis of spectra collected for Ir(III), OH
ꢀ
ꢀ
ꢀ
and IO3 solutions and Ir(III), OH , IO and maltose
3
solutions. From the observed spectra it is apparent that
ꢀ
3
when iodate solutions of concentrations 1 · 10 and
ꢀ
3
2
· 10 M were added to the solution of Ir(III) chloride
ꢀ
and OH , an increase in absorbance from 2.08 to 2.54
and 2.62 was obtained (Fig. 7: (3), (5) and (6)). This
spectral information led us to conclude that the above
equilibrium exists in solution, indicating an increasing
2
–
O
O
formation of the complex,
.
ꢀ
ꢀ
[
IrCl (H O) OH] and IO3 , spectra for the solution
I — O–IrCl (H O)OH
3
2
2
3
2
ꢀ
of Ir(III) chloride and OH and for the solutions of Ir-
III) chloride and OH with two different concentrations
When maltose solutions of two different concentrations
ꢀ
ꢀ
ꢀ
(
were added to the solution of Ir(III), OH and IO3
again an increase in absorbance was noted (Fig. 7: (5),
,
ꢀ
of IO3 have been collected (Fig. 7: (3), (5) and (6)).
From the spectra, it is clear that with the addition of
potassium iodate solution there is an increase in absor-
bance from 2.08 to 2.54 and 2.62, with a slight shift in
kmax value towards longer wavelength, that is, from
(
(
7) and (8)). This shows that the complex formed
kmax 217 nm) subsequently combines with the reac-
tive form of maltose to give a most reactive activated
2-
OH
R — C
O
2
16 to 217 and 220 nm. This increase in absorbance with
ꢀ
C — H
O
the increase in ½IO3 ꢁ can be considered as an indication
of increased formation of the complex between a reac-
tive species of Ir(III) chloride and a reactive species of
potassium iodate in alkaline medium according to the
following equilibrium:
complex of the type
according
O
I
O – IrCl3(H2O)2
2
–
O
O
-
-
[
IrCl (H O ) O H]
+
IO
H O
2
3
2
2
3
I — O –IrCl (H O )O H
+
3
2
The shift in kmax value towards longer wavelength is
due to the combination of a chromophore, IO3 , and an
to the following equilibrium:
ꢀ
2
-
2
-
OH
R — C
O
O
O
H − C − OH
C — H
O
|
|
I
O
IrCl3(H2O)OH
+
C − OH
O
|
I
R
O – IrCl (H O)
3 2 2

A. K. Singh et al. / Carbohydrate Research 342 (2007) 1078–1090
1085
4.5. Reaction mechanisms for the oxidation of xylose and
maltose
C
2
¼ K
1
½C
1
ꢁ½OHꢀꢁ
ð7Þ
and
C ¼ K ½C ꢁ½IO ꢀꢁ
Eqs. 7 and 8 will give Eq. 9
ð8Þ
On the basis of (1) the discussion made above for the
reactive species of Ir(III) chloride, potassium iodate,
and also for the reactive form of reducing sugars, (2)
spectral information collected for the formation of a
complex or complexes during the course of reaction
3
2
2
3
ꢀ
C
3
¼ K
1
K
2
½C
1
ꢁ½OH ꢁ½IO ꢀꢁ
ð9Þ
3
On substituting the value of C from Eq. 9 to Eq. 6, we
3
ꢀ
and (3) kinetic orders observed with respect to ½IO ꢁ,
3
get Eq. 10
ꢀ
[
Ir(III)], [OH ], [xylose] and [maltose], the reactions
d½IO ꢀꢁ
shown in Schemes 1 and 2 can be proposed for the oxi-
dation of xylose and maltose, respectively.
3
ꢁ½OH ꢁ½IO ꢀꢁ
ꢀ
rate ¼ ꢀ
¼ 2k
3
K
1
K
2
½C
1
ð10Þ
3
dt
On the basis of the above reaction scheme, the rate in
terms of decrease in concentration of IO3 can be
expressed as
According to Scheme 1, the total concentration of
ꢀ
Ir(III), that is, [Ir(III)]
be shown as Eq. 11
at any time in the reaction can
T
½
IrðIIIÞꢁ ¼ ½C
1
ꢁ þ ½C
2
ꢁ þ ½C
3
ꢁ
ð11Þ
d½IO3ꢀꢁ
T
rate ¼ ꢀ
¼ 2k
3
½C
3
ꢁ;
ð6Þ
dt
Eq. 12 based on Eqs. 7, 9 and 11 is as follows
½
IrðIIIÞꢁ
T
2
where 2 indicates that 1 mol of xylose is oxidized by
½C
1
ꢁ ¼ 1
ꢀ
ꢀ
ð12Þ
þ K
1
½OH ꢁ þ K
1
K
½OH ꢁ½IO3ꢀꢁ
2
mol of KIO3.
On applying the law of chemical equilibrium to steps
On substituting the value of [C ] from Eq. 12 to Eq. 10,
1
(
i) and (ii), Eqs. 7 and 8, respectively, are obtained:
we get Eq. 13
K
1
−
−
[
IrCl (H O) ] + OH
[IrCl (H O) OH] + H O
(i)
3
2
3
3
2
2
2
(
C )
(C2)
1
2-
O
O
K
2
−
−
[
IrCl (H O) OH] +IO
(ii)
3
2
2
3
I
O
IrCl3(H2O)OH
+ H O
2
(
C )
(C3)
2
2
-
2-
O
O
O
k
3
(iii)
I
O
IrCl3(H2O)OH
O
I
IrCl3(OH)
+ H O
2
slow and rate
determining step
O
(
C )
3
(
C )
4
2-
−
H − C − OH
O
|
O
||
O
|
|
–
fast
−
2
2−
O
I
IrCl3(OH)
+
C − OH
+ 2OH ⎯⎯→ IO
+
(iv)
R – C — C+[IrCl (H O)(OH) ]
|
3
2
2
O
|
|
H
R
OH
(
C )
4
(S)
where R stand for C H O
3
3
7
2
–
–
[IrCl
3
(H
2
O)(OH)
2
]
+ 2H
2
O
[IrCl
3
(H
2
O)
3
]
+
2OH
(v)
−
O
|
O
||
–
–
IO / OH / Ir(III)
3
−
R – C — C
HCOOH + RCOOH + IO2 +IrCl (H O)
Formic acid Threonic acid
3
(vi)
3
2
|
|
OH
H
Scheme 1.

1086
A. K. Singh et al. / Carbohydrate Research 342 (2007) 1078–1090
K
1
−
−
[
IrCl (H O) ] + OH
[IrCl (H O) OH] + H O
(i)
3
2
3
3
2
2
2
(
C )
(C2)
1
2-
O
O
K
2
−
−
[
IrCl (H O) OH] +IO
I
O
IrCl (H O)OH
+ H O (ii)
2
3
2
2
3
3
2
(
C )
2
(C3)
2
-
2
-
OH
O
O
O
O
H − C − OH
|
|
K
3
R — C
O
C — H
O
(
iii)
I
IrCl3(H2O)OH
+
C − OH
|
I
R
O – IrCl (H O)
3
2
2
(
C )
(S)
(C4)
3
where R stands for C H O
4
9
4
2
-
−
OH
R — C
O
O
|
O
||
C — H
O
k
4
−
+
H
2
O
IO + R – C — C + [IrCl
3
(H
O)
2 3
]
(iv)
O
slow and rate
determining step
2
|
|
I
OH
H
O – IrCl (H O)
3
2
2
(
C )
4
−
O
|
O
||
–
–
IO / OH / Ir(III)
3
−
R – C — C
HCOOH + RCOOH + IO + IrCl (H O)
3
(v)
2
3
2
|
|
fast
Formic acid Arabinonic acid
OH
H
Scheme 2.
d½IO3ꢀꢁ
dt
d½IO ꢀꢁ 2k
K
K
½OH ꢁ½IO ꢁ½IrðIIIÞꢁT
ꢀ
ꢀ
3
3
1
2
3
rate ¼ ꢀ
rate ¼ ꢀ
¼
ð16Þ
ꢀ
dt
1 þ K
1
½OH ꢁ
ꢀ
ꢀ
2
k
3
K
1
K
2
½OH ꢁ½IO ꢁ½IrðIIIÞꢁ
3
T
¼ 1
ð13Þ
ꢀ
ꢀ
ꢀ
If it is assumed that K ½IO ꢁ ꢂ 1 at high concentra-
2
3
þ K
1
½OH ꢁ þ K
1
K
2
½OH ꢁ½IO3ꢀꢁ
ꢀ
tions of IO , then under this condition, 1 will be negli-
3
gible in comparison to K ½IO ꢀꢁ and Eq. 14 will be
The rate law 13 derived on the basis of proposed Scheme
2
3
1
is in close agreement with the kinetic information
reduced to Eq. 17
ꢀ
obtained experimentally about the effect of ½IO ꢁ,
d½IO ꢀꢁ 2k K K ½OH ꢁ½IO ꢁ½IrðIIIÞꢁ
ꢀ
ꢀ
3
3
3
1
2
3
ꢀ
T
ꢀ
rate ¼ ꢀ
¼
[
Ir(III)] , [OH ] and [S] on the rate of oxidation of
T
ꢀ
3
dt
1 þ K K ½OH ꢁ½IO
ꢁ
1
2
xylose.
ð17Þ
ð18Þ
Eq. 13 can also be written as
According to Eq. 7
d½IO3ꢀꢁ 2k
K
1 þ K
K
½OH ꢁ½IO ꢁ½IrðIIIÞꢁ
ꢀ
ꢀ
3
1
2
3
T
rate ¼ ꢀ
¼
ð14Þ
ð15Þ
ꢀ
½C ꢁ
ꢀ
2
dt
1
½OH ꢁf1 þ K
2
½IO3 ꢁg
ꢀ
K ½OH ꢁ ¼
1
½
C
1
ꢁ
According to Eq. 8
Equilibrium reaction (i) shows that, at low concentra-
tions of hydroxyl ion, [C ] will be less than [C ], and
½
½
C
C
3
2
ꢁ
ꢁ
K ½IO3ꢀꢁ ¼
2
1
2
ꢀ
hence the inequality K [OH ] ꢅ 1 can be assumed as
1
valid one which will lead Eq. 16 to take the form of
Eq. 19.
ꢀ
Since at low concentration of IO3 , [C ] ꢅ [C ], hence
3
2
ꢀ
according to Eq. 15 the inequality K
2
½IO ꢁ ꢅ 1 can
3
d½IO ꢀꢁ
be assumed as a valid one. Under this condition,
Eq. 14 will become Eq. 16
3
½OH ꢁ½IO ꢀꢁ½IrðIIIÞꢁ
ꢀ
rate ¼ ꢀ
¼ 2k
3
K
1
K
2
3
T
dt

A. K. Singh et al. / Carbohydrate Research 342 (2007) 1078–1090
1087
or
Table 3. The calculated values of composite rate constants for the
ꢀ
ꢀ
variations of ½IO ꢁ and [OH ] in the Ir(III)-catalyzed oxidation of
3
d½IO3ꢀꢁ
dt
xylose and maltose at 40 ꢁC
¼ k ½OH ꢁ½IO ꢀꢁ½IrðIIIÞꢁ
0
ꢀ
ð19Þ
rate ¼ ꢀ
3
T
ꢀ
3
ꢀ
2
½IO3 ꢁ ꢃ 10
[OH ] · 10
Composite rate constant
(
mol/L)
(mol/L)
0
2
00
4
k · 10
k · 10
0
where k = 2k K K .
ꢀ2
2
ꢀ1
ꢀ3
3
ꢀ1
)
3
1 2
(
mol L s
)
(mol L s
Eq. 19 is the rate law, valid for low concentrations of
OH and IO3 and for all concentrations of Ir(III) lying
Xylose
Maltose
ꢀ
ꢀ
0.4
.8
1.0
1.2
1.6
20.00
20.00
20.00
20.00
20.00
20.00
20.00
5.00
10.00
15.00
20.00
25.00
30.00
35.00
1.11
1.16
1.12
1.20
1.17
1.19
1.17
1.07
1.12
1.12
1.12
1.12
1.09
—
0.39
0.36
0.37
0.37
0.36
0.37
0.35
—
0.38
0.37
0.39
0.37
0.37
0.36
ꢀ
6
ꢀ6
between 6.69 · 10 M and 53.52 · 10 M and is well
0
in agreement with experimental findings.
At high concentrations of IO3
ꢀ
ꢀ
and OH ,
ꢀ
ꢀ
K
1
K
2
½IO3 ꢁ½OH ꢁ will be very much greater than 1,
2
2
1
1
1
.0
.4
.00
.00
.00
and under this condition, Eq. 17 will finally be converted
into Eq. 20
d½IO3ꢀꢁ
rate ¼ ꢀ
¼ 2k ½IrðIIIÞꢁ
ð20Þ
3
T
dt
1.00
1
1
1
.00
.00
.00
ꢀ
Eq. 20 is valid only for higher concentrations of IO
3
ꢀ
and OH . Observed zero-order kinetics with respect to
each ½IO3 ꢁ and [OH ] at their higher concentrations
ꢀ
ꢀ
ꢀ2
Solution conditions: [Sugar] = 2.00 · 10 mol/L, [Ir(III)] = 26.76 ·
are well in accordance with the rate law 20.
ꢀ6
ꢀ6
1
0
mol/L (xylose) and 13.38 · 10 mol/L (maltose), [l] = 0.25 mol/
ꢀ
ꢀ
Using Eq. 19, and making use of the kinetic data ob-
L (for ½IO3 ꢁ variation) and 0.55 mol/L (for [OH ] variation).
ꢀ
ꢀ
tained for low ½IO3 ꢁ and [OH ] at 40 ꢁC and under uni-
form conditions, the values of a composite rate constant
k ) have been calculated and are presented in Table 3.
Almost the same values of k for the variations of
0
With the help of Eqs. 22–24 and 26, we can write Eq. 27
(
0
½IrðIIIÞꢁ
T
ꢀ
½
C
1
ꢁ ¼
ꢀ
ꢀ
ꢀ
ꢀ
½OH ꢁ½IO ꢀꢁ þ K
½IO3ꢀꢁ½OH ꢁ½Sꢁ
ð27Þ
½
IO3 ꢁ and [OH ] in their lower range, clearly prove
1 þ K
1
½OH ꢁ þ K
1
K
2
1 2 3
K K
3
the validity of the rate law 19 and in turn the rate law
1
3 and hence the proposed mechanism.
On substituting the value of [C ] from Eq. 27 to Eq. 25,
On the basis of above reaction Scheme 2 and stoichio-
1
metric data, the rate in terms of disappearance of ½IO ꢀꢁ
we get Eq. 28
3
d½IO ꢀꢁ
can be written as Eq. 21
3
rate ¼ ꢀ
dt
d½IO3ꢀꢁ
ꢀ
ꢀ
4k K K K ½OH ꢁ½IO ꢁ½Sꢁ½IrðIIIÞꢁ
4
1
2
3
rate ¼ ꢀ
¼ 4k ½C ꢁ
ð21Þ
3
T
4
4
¼ 1
ꢀ
ꢀ
ꢀ
½OH ꢁ½IO3ꢀꢁ½Sꢁ
ð28Þ
ꢀ
dt
þ K
1
½OH ꢁ þ K
1
K
2
½OH ꢁ½IO ꢁ þ K
1 2 3
K K
3
where 4 indicates that 1 mol of maltose is oxidized by
mol of KIO3.
On applying the law of chemical equilibrium to steps
i), (ii) and (iii), we get Eqs. 22–24, respectively.
4
Eq. 28 can also be written as Eq. 29
d½IO ꢀꢁ
(
3
rate ¼ ꢀ
dt
ꢀ
ꢀ
ꢀ
C
2
¼ K
1
½C
1
ꢁ½OH ꢁ
ð22Þ
ð23Þ
4
k
4
K
1
K
2
K
3
½OH ꢁ½IO ꢁ½Sꢁ½IrðIIIÞꢁ
3
ꢀ
T
¼
C ¼ K K ½C ꢁ½OH ꢁ½IO ꢀꢁ
ꢀ
ꢀ
ꢀ
3
1
2
1
3
1 þ K ½OH ꢁ þ K K ½OH ꢁ½IO ꢁf1 þ K ½Sꢁg
1
1
2
3
3
ð29Þ
and
At low concentrations of S, the inequality K [S] ꢅ 1 can
3
ꢀ
ꢁ½OH ꢁ½IO3ꢀꢁ½Sꢁ
C
4
¼ K
1
K
2
K
3
½C
1
ð24Þ
be assumed as valid and under this condition Eq. 29 will
be reduced to Eq. 30
On putting the value of C from Eq. 24 into Eq. 21, we
4
d½IO ꢀꢁ
have Eq. 25
3
rate ¼ ꢀ
d½IO3ꢀꢁ
dt
ꢀ
ꢀ
ꢀ
ꢀ
rate ¼ ꢀ
¼ 4k
4
K
1
K
2
K
3
½C
1
ꢁ½OH ꢁ½IO3ꢀꢁ½Sꢁ ð25Þ
4k K K K ½IO ꢁ½OH ꢁ½Sꢁ½IrðIIIÞꢁ
4
1
2
3
3
T
dt
¼
ð30Þ
ꢀ
½IO ꢁg
3
1
þ K
1
½OH ꢁf1 þ K
2
According to mechanism, at any moment in the reaction
the total concentration of Ir(III), that is, [Ir(III)] can be
expressed as Eq. 26
At high concentrations of S, Eq. 29 will be reduced to
Eq. 31 because in this range, the existence of the inequal-
T
ity K [S] ꢂ 1 will make 1 negligible in comparison to
3
½
IrðIIIÞꢁ ¼ ½C
1
ꢁ þ ½C
2
ꢁ þ ½C
3
ꢁ þ ½C
4
ꢁ
ð26Þ
T
K [S].
3

1088
A. K. Singh et al. / Carbohydrate Research 342 (2007) 1078–1090
d½IO3ꢀꢁ
dt
ꢀ
ꢀ
OH] and IO , as a result of which the transition state
3
rate ¼ ꢀ
will be a more highly charged species. Due to this more
solvent molecules will be required than for the separate
ions. This would lead to a decrease in entropy. The
observed negative entropy of activation supports
the formation of aforesaid activated complex by the
interaction of two similarly charged species in step (ii)
of Scheme 1. For the oxidation of maltose, a separate
reaction path (Scheme 2) has been proposed,
ꢀ
ꢀ
4
k K K K ½IO ꢁ½OH ꢁ½Sꢁ½IrðIIIÞꢁ
4
1
2
3
3
T
¼
ð31Þ
ꢀ
ꢀ
1
þ K ½OH ꢁf1 þ K K ½IO ꢁ½Sꢁg
1
2
3
3
ꢀ
Further, when at low concentrations of IO3
;
ꢀ
ꢀ
K ½IO ꢁ ꢅ 1 then K ½IO ꢁ in comparison to 1 can be
2
3
2
3
neglected in Eq. 30, which under this condition will be
reduced to Eq. 32.
d½IO3ꢀꢁ
where the most reactive activated complex,
rate ¼ ꢀ
dt
2-
OH
R — C
O
ꢀ
ꢀ
4
k
4
K
1
K
2
K
3
½IO3 ꢁ½OH ꢁ½Sꢁ½IrðIIIÞꢁ
C — H
O
T
¼
ð32Þ
ꢀ
has been shown to be formed by
1
þ K
1
½OH ꢁ
O
I
ꢀ
At high concentrations of IO3 and S, the inequality
O – IrCl (H O)
3 2 2
K
2
K
3
½IO3ꢀꢁ½Sꢁ ꢂ 1 will reduce Eq. 31 to Eq. 33.
ꢀ
the
interaction
of
a
charged
species
d½IO3
ꢁ
rate ¼ ꢀ
2-
dt
O
O
ꢀ
ꢀ
and
a
neutral molecule
4
k
4
K
1
K
2
K
3
½IO ꢁ½OH ꢁ½Sꢁ½IrðIIIÞꢁ
3
T
I
O
3 2
IrCl (H O)OH
¼
ð33Þ
ꢀ
ꢀ
1
þ K
1
K
2
K
3
½IO3 ꢁ½Sꢁ½OH ꢁ
H—C—OH
From Eqs. 32 and 33, it is clear that at low concentra-
ꢀ
ꢀ
. In this case, the transition state will have the
tions of OH , Eq. 32 with inequality K [OH ] ꢂ 1,
C—OH
R
1
ꢀ
and at high concentrations of OH , Eq. 33 with inequal-
ꢀ
ꢀ
ity K
3
1
K
2
K
3
½IO3 ꢁ½Sꢁ½OH ꢁ ꢂ 1, will be reduced to Eqs.
same charge, but it will be dispersed over a greater vol-
ume, as a result of which the transition state will be less
polar than the reactants. This would lead to an increase
in entropy. The observed positive entropy of activation
in the Ir(III)-catalyzed oxidation of maltose is certainly
evidence for step (iii) of Scheme 2. In step (iii) the com-
plex in the transition state has the same charge dispersed
over greater volume than the charged species while inter-
acting with the neutral species in the reactant state.
Since the order of the frequency factor in the oxidation
of xylose and maltose is not the same, it is reasonable to
assume that the proposed reaction mechanisms as
shown in Schemes 1 and 2 are well in accordance with
the observed kinetic data.
4 and 35, respectively.
d½IO3ꢀꢁ
dt
rate ¼ ꢀ
4k
½IO3ꢀꢁ½OH ꢁ½Sꢁ½IrðIIIÞꢁ
ꢀ
ð34Þ
ð35Þ
¼
4
K
1
K
2
K
3
T
and
d½IO3ꢀꢁ
rate ¼ ꢀ
¼ 4k ½IrðIIIÞꢁ
4
T
dt
The rate law 34 can also be written as
d½IO3ꢀꢁ
00
ꢀ
rate ¼ ꢀ
¼ k ½IO3ꢀꢁ½OH ꢁ½IrðIIIÞꢁ ½Sꢁ
ð36Þ
T
dt
00
where k = 4k K K K .
4
1 2 3
When the validity of the rate law 34 or 36 was tested
by the kinetic data observed at low concentrations of
substrate under uniform experimental conditions for
4
.6. Comparative studies
ꢀ
ꢀ
The results of the present study of the Ir(III)-catalyzed
oxidation of xylose and maltose by potassium iodate
in alkaline medium were compared with the results
reported for Ru(III)-catalyzed oxidation of indigo car-
mine by iodate in acidic medium and the Ir(III)-
catalyzed oxidation of reducing sugars by NBA in
acidic medium. When the present study in respect of
the role of iodate was compared with the Ru(III)-cata-
lyzed oxidation of indigo carmine, it is found that
the species IO3 is the reactive species of potassium
iodate in acidic as well as in alkaline medium. As far
as the kinetic order with respect to IO3 is concerned,
the variations of OH and IO3 , it was found that the
00
values of k remain almost the same for both the varia-
tions in their lower range (Table 3). This proves the
validity of rate law 34 and in turn rate law 29, and hence
the proposed mechanism.
2
4
6
In the present study of the Ir(III)-catalyzed oxida-
tion of reducing sugars, Schemes 1 and 2 have been
proposed for xylose and maltose, respectively. In
Scheme 1, the most unstable activated complex,
2
4
ꢀ
2-
O
O
,
is formed by the interaction
ꢀ
I
O
IrCl (H O)OH
3 2
ꢀ
it is first-order throughout the variation of ½IO ꢁ in
3
2
4
of two similarly charged species, that is, [IrCl (H O) -
the reported Ru(III)-catalyzed oxidation and first to
3
2
2

A. K. Singh et al. / Carbohydrate Research 342 (2007) 1078–1090
1089
zero order in the present study of the Ir(III)-catalyzed
2-
O
O
oxidation of reducing sugars. When the present study
has been made for the effect of [sugar] on the rate of oxi-
dation, it is found that in the case of xylose the rate is
independent of [xylose] and in the case of maltose, it is
first order at low concentrations of maltose and becomes
(
2) The complex,
and the com-
I
O
3 2
IrCl (H O)OH
2
-
OH
O
R — C
C — H
O
plex,
have been proposed as
6
O
zero order at its higher concentrations. The reported
I
zero-order kinetics in [sugar] for Ir(III)-catalyzed oxida-
tion seems to be similar with the present study as far as
order with respect to xylose is concerned, but dissimilar
as far as the order with respect to maltose is concerned.
Since the present study has been performed in alkaline
medium, the reducing sugar molecule participates in
the reaction in the enediol form, whereas in the
O – IrCl (H O)
3
2
2
the most unstable activated complex in the oxida-
tion of xylose and maltose, respectively.
3) Step (ii) of Scheme 1, in which the interaction
between two similarly charged species results in
the formation of an activated complex, is supported
by observed negative entropy of activation. Step
(
6
reported Ir(III)-catalyzed oxidation, it participates in
the reaction as such. On the basis of the observed kinetic
(
iii) of Scheme 2, in which the interaction between
a charged species and a natural molecule results
in the formation of an activated complex, is sup-
ported by the observed positive entropy of
activation.
data and spectral information, it is concluded that
ꢀ
[
IrCl (H O) OH] is the reactive species of Ir(III) chlo-
3
2
2
ride in the present study of the oxidation of reducing
sugars in aqueous alkaline medium, whereas in the
6
(4) First-order kinetics observed at low concentrations
of maltose distinguishes Scheme 2 from Scheme 1
where zero-order kinetics with respect to [xylose]
throughout its variation has been observed.
reported Ir(III)-catalyzed oxidation of reducing sugars,
the reactive species of Ir(III) chloride was found to be
3
ꢀ
[
IrCl6] . The proposed reaction Schemes 1 and 2 in
the present study clearly show that there is formation
of a complex between the reactive species of Ir(III) chlo-
ride and reactive species of potassium iodate in alkaline
medium, which is well supported by the observed kinetic
data and also by the spectral evidence collected. This
finding in the present paper is clearly in line with the
(
5) The rate of oxidation of either xylose or maltose is
unaffected by the ionic strength of the medium.
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1
1