Oxidation of butane-1,3-, butane-1,4-, 2-methyl pentane-2,4- and 3-methyl pentane-2,4-diols by cerium(IV) in aqueous acidic medium catalyzed by rhodium(III)

Article abstract of DOI:10.1007/s11243-011-9527-6

The oxidation kinetics of butane-1,3-diol, butane-1,4-diol, 2-methyl pentane-2,4-diol and 3-methyl pentane-2,4-diol with cerium(IV) catalyzed by rhodium(III) in aqueous sulfuric acid showed a peculiar nature with respect to the variation in oxidant concentration, such that the reaction follows first-order kinetics in [Ce(IV)] at low [Ce(IV)] and then reaches a maximum with increasing [Ce(IV)], beyond which further increase in the oxidant concentration retards the rate. The rate shows direct proportionality with respect to [diol] at low concentrations, becoming independent of [diol] at higher concentrations. The rate is first order in catalyst. Retarding effects are observed when [H ⁺] and [Ce(III)] are increased, while [Cl^-] and hence ionic strength have a positive effect on the rate. Spectroscopic studies confirmed that the primary hydroxyl groups in butane-1,3-diol and butane-1,4-diol resulted in the formation of 3-hydroxy butanal and 4-hydroxy butanal, respectively. In the case of oxidation of the secondary hydroxyl groups in 2-methyl pentane-2,4-diol and 3-methyl pentane-2,4-diol, the products of oxidation were 4-hydroxy-4-methyl pentan-2-one and 4-hydroxy-3-methyl pentan-2-one, respectively.

Full text of DOI:10.1007/s11243-011-9527-6

Transition Met Chem (2011) 36:739–746
DOI 10.1007/s11243-011-9527-6
Oxidation of butane-1,3-, butane-1,4-, 2-methyl pentane-2,4-
and 3-methyl pentane-2,4-diols by cerium(IV) in aqueous acidic
medium catalyzed by rhodium(III)
•
Praveen K. Tandon Shaista Z. Khanam
•
•
Suresh C. Yadav Ritesh C. Shukla
Received: 29 April 2011 / Accepted: 8 August 2011 / Published online: 28 August 2011
Ó Springer Science+Business Media B.V. 2011
Abstract The oxidation kinetics of butane-1,3-diol,
butane-1,4-diol, 2-methyl pentane-2,4-diol and 3-methyl
pentane-2,4-diol with cerium(IV) catalyzed by rhodium(III)
in aqueous sulfuric acid showed a peculiar nature with
respect to the variation in oxidant concentration, such that
the reaction follows first-order kinetics in [Ce(IV)] at low
[Ce(IV)] and then reaches a maximum with increasing
[Ce(IV)], beyond which further increase in the oxidant
concentration retards the rate. The rate shows direct pro-
portionality with respect to [diol] at low concentrations,
becoming independent of [diol]at higher concentrations. The
rate is first order in catalyst. Retarding effects are observed
when [H?] and [Ce(III)] are increased, while [Cl-] and hence
ionic strength have a positive effect on the rate. Spectro-
scopic studies confirmed that the primary hydroxyl groups in
butane-1,3-diol and butane-1,4-diol resulted in the formation
of 3-hydroxy butanal and 4-hydroxy butanal, respectively.
In the case of oxidation of the secondary hydroxyl groups
in 2-methyl pentane-2,4-diol and 3-methyl pentane-2,4-diol,
the products of oxidation were 4-hydroxy-4-methyl pentan-
2-one and 4-hydroxy-3-methyl pentan-2-one, respectively.
point of view. However, catalysis of these oxidations by
transition metal ions has been less well studied. We have
reported a kinetic study of the catalyzed oxidation of var-
ious aliphatic organic compounds such as cyclic ketones by
cerium(IV) [1]. We have also reported on the catalytic
activity of rhodium(III) in the oxidation of cyclic ketones
by alkaline hexacyanoferrate(III), in which it was observed
that the catalytically active species of rhodium(III) depends
on the pH of the medium [2, 3]. Rhodium(III) proved to be
more effective than iridium(III) and palladium(II) in cata-
lyzing organic oxidations with hydrogen peroxide from a
synthetic point of view [4]. In the present work, we have
studied the oxidation of four diols, namely butane-1,3-diol,
butane-1,4-diol, 2-methyl pentane-2,4-diol and 3-methyl
pentane-2,4-diol by cerium(IV) sulfate in aqueous acidic
medium catalyzed by rhodium(III) chloride. These diols
were chosen in order to investigate the potential of rho-
dium(III) in catalyzing the oxidation of non-vicinal diols;
to confirm whether a C–H or C–C bond is broken during
the oxidation; to see whether the methyl group instead of
hydrogen has any effect on the oxidation of the hydroxyl
group; and to confirm whether the secondary or tertiary
alcoholic group is oxidized under these conditions.
Introduction
The uncatalyzed oxidation of aliphatic compounds by
cerium(IV) has been frequently reported from the kinetic
Results and discussion
For all the four diols, initially, the first-order rate constant
k values are constant but decrease at higher [Ce(IV)]
(Table 1). The rate values (-dc/dt) initially increase with
increasing [Ce(IV)] and then reach a maximum beyond
which further increase in [Ce(IV)] decreases the rate
(Fig. 1). Plots of rate versus [diol] for all substrates gave a
straight line passing through the origin at low [diol], but at
higher organic substrate concentrations, the trend of the
Electronic supplementary material The online version of this
article (doi:10.1007/s11243-011-9527-6) contains supplementary
material, which is available to authorized users.
P. K. Tandon (&) ꢀ S. Z. Khanam ꢀ S. C. Yadav ꢀ R. C. Shukla
Department of Chemistry, University of Allahabad,
Allahabad 211002, India
e-mail: pktandon123@rediffmail.com
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Transition Met Chem (2011) 36:739–746
Table 1 Effect of variation in [Cerium(IV)] and [Diols] on the rate at 32 °C
[CeIV] 9 104 M
k 9 102 min-1
[Diol] 9 103 M
k* (M-1 min-1
)
A
B
C
D
A
B
C
D
2.00
3.00
4.00
5.00
6.00
7.00
8.00
10.0
2.80
2.67
2.92
2.84
2.78
2.14
1.66
1.00
0.50
0.47
0.40
0.46
0.45
0.47
0.31
0.14
0.50
0.57
0.50
0.50
0.50
0.31
0.25
0.15
0.32
0.33
0.36
0.33
0.33
0.21
0.17
0.13
2.00
2.50
3.00
4.00
5.00
6.00
8.00
10.0
4.20
4.32
4.26
4.16
4.33
4.44
4.00
3.50
0.96
0.96
0.89
0.83
0.72
0.66
0.55
0.50
0.86
0.80
0.85
0.80
0.73
0.71
0.67
0.64
0.86
–
0.89
0.90
0.80
0.83
0.81
0.80
A-butane-1,3-diol; B-butane-1,4-diol; C-2-methyl pentane-2,4-diol; D-3-methyl pentane-2,4-diol
Variation in cerium(IV) sulfate concentrations: [Diol] 9 10-3 M = 5.0 (A to D); [H2SO4] 9 M = 0.50 (A, D), 1.00 (B), 0.35 (C);
[RhCl3] 9 10-5 M = 2.39 (A, B), 1.40 (C, D)
Variation in diol concentrations: [Ce(SO4)2] 9 10-4 = 5.00 M (A to D), [H2SO4] 9 M = 0.50 (A), 1.00 (B), 0.20 (C), 0.25 (D).
[RhCl3] 9 10-5 M = 2.39 (A, B), 1.40 (C and D)
k* = -dc/dt/[Ce(IV)][diol]
[Ce(SO4)2] x 103 M
12
10
8
6
4
2
0
3.5
3
18
15
12
9
2.5
2
C
1.5
1
A
B
6
D
3
0.5
0
0
0
2
4
6
8
10
12
[Ce(SO4)2] x 103 M
Fig. 1 Effect of variation in [Ce(SO4)2] on the rate at 32 °C.
[Diol] 9 10-3 M = 5.0 (A–D); [H2SO4] 9 M = 0.50 (A, D), 1.00
(B), 0.35 (C); [RhCl3] 9 10-5 M = 2.39 (A, B), 1.40 (C, D).
A-butane-1,3-diol; B-butane-1,4-diol; C-2-methyl pentane-2,4-diol;
D-3-methyl pentane-2,4-diol
Fig. 2 Effect of variation in [Diol] on the rate at 32 °C.
[Ce(SO4)2] 9 10-4 = 5.00 M (A–D), [H2SO4] 9 M = 0.50 (A),
1.00 (B), 0.20 (C), 0.25 (D). [RhCl3] 9 10-5 M = 2.39 (A, B),
1.40 (C and D). A-butane-1,3-diol; B-butane-1,4-diol; C-2-methyl
pentane-2,4-diol; D-3-methyl pentane-2,4-diol
line to become parallel to the x-axis becomes apparent
(Fig. 2). This is confirmed by the trend in the second-order
rate constant k* values (Table 1). The rate increases pro-
portionally with increasing concentrations of rhodium(III)
for all substrates. The second-order rate constant k* values
were determined as 1.16 0.13, 1.67 0.05, 1.65 0.25
and 2.22 0.23 for butane-1, 3-diol, butane-1, 4-diol, 2-
methyl pentane-2,4-diol and 3-methyl pentane-2,4-diol,
respectively (Table 2). Straight lines passing through the
origin were obtained on plotting -dc/dt versus [RhCl3],
indicating first-order kinetics in [RhCl3]. This was further
confirmed from plots of log (-dc/dt) versus log [RhCl3],
which gave slopes of 0.96, 0.98, 0.97 and 0.98 for butane-
1, 3-diol, butane-1, 4-diol, 2-methyl pentane-2,4-diol and
3-methyl pentane-2,4-diol, respectively. The rate of reac-
tion decreases with increasing concentrations of sulfuric
acid (Table 2) and also on addition of Ce2(SO4)3 to the
reaction mixture (Table 3). On plotting rate values against
[H2SO4] or [Ce2(SO4)3], an initial sharp fall in the rate
becomes less prominent at higher concentrations of both
acid and cerium(III). These observations suggest that the
reduction of cerium(IV) to cerium(III) along with the
removal of H? takes place before the rate determining step
in the mechanism. Where the dielectric constant of the
medium was increased by addition of a standard solution
of potassium chloride, the rate was found to increase
(Table 3).
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Transition Met Chem (2011) 36:739–746
741
Table 2 Effect of variation in [RhCl3] and [H?] on the rate at 32 °C
[RhCl3] x 105 M
k* 9 10-2 (M-1 min-1
)
[H2SO4] 9 M
-dc/dt 9 106 (M min-1
)
A
B
C
D
A
B
C
D
0.47
0.75
0.95
1.70
1.90
2.39
3.10
3.58
4.70
–
1.44
–
–
–
2.68
–
0.30
0.50
0.60
0.70
0.80
1.00
1.20
1.30
1.50
1.60
1.75
1.85
2.00
–
8.28
3.96
–
19.80
6.66
4.98
–
1.60
1.68
1.76
–
1.78
2.09
–
–
13.20
1.17
–
2.09
–
–
12.48
2.49
–
3.96
–
–
3.00
–
1.05
1.11
1.13
1.11
–
1.40
1.46
1.61
1.58
–
2.11
2.09
2.15
2.23
–
–
1.67
1.61
1.68
1.70
–
1.99
–
2.82
2.46
–
9.00
2.49
–
–
–
–
1.26
0.71
–
–
1.50
–
8.28
–
–
–
–
1.65
1.39
1.39
1.29
–
–
–
–
–
–
–
–
–
–
–
–
–
–
–
–
–
–
–
–
–
–
–
–
A-butane-1,3-diol; B-butane-1,4-diol; C-2-methyl pentane-2,4-diol; D-3-methyl pentane-2,4-diol
Variation in RhCl3 concentrations: [Ce(SO4)2] 9 10-4 M = 5.00 (A to D); [Diol] 9 10-3 M = 5.0 (A to D); [H2SO4] 9 M = 0.30 (A), 1.0 M
(B), 0.30 (C), 0.50 (D)
Variation in H2SO4 concentrations: [Ce(SO4)2] 9 10-4 M = 5.00 (A to B); [Diol] 9 10-3 M = 5.0 (A to B); [RhCl3] 9 10-5 M = 2.39 (A, B),
1.40 (C, D)
k* = -dc/dt/[Ce(IV)][diol]
Table 3 Effect of variation in [CeIII] and [Cl-] on the rate at 32 °C
[Ce2(SO4)3] 9 104 M
-dc/dt 9 106 (M min-1
)
[KCl]
M
-dc/dt 9 106 (M min-1
)
A
B
C
D
A
B
C
D
2.00
2.50
3.50
5.00
7.00
7.50
10.00
1.67
1.00
1.00
0.75
–
1.17
1.00
0.75
0.67
0.44
–
8.33
7.50
7.50
6.67
–
1.50
–
0.70
0.85
1.00
1.15
1.30
1.50
–
2.75
4.00
4.30
–
1.67
3.00
4.38
5.00
5.00
6.25
–
6.25
0.71
–
7.50
7.50
8.30
10.00
12.50
–
1.33
1.00
–
1.00
–
4.50
5.00
–
1.50
1.75
–
0.60
0.33
5.00
4.00
0.71
–
0.33
A-butane-1,3-diol; B-butane-1,4-diol; C-2-methyl pentane-2,4-diol; D-3-methyl pentane-2,4-diol
Variation in Ce2(SO4)3 concentrations: [Ce(SO4)2] 9 10-4 M = 5.00 (A to D); [Diol] 9 10-3 M = 5.0 (A to B); [RhCl3] 9 10-5 M = 2.39 (A,
B), 1.40 (C and D); [H2SO4] 9 M = 0.30 (A); 0.50 (B), 0.30 (C), 0.70 (D)
Variation in KCl concentrations: A-[Ce(SO4)2] 9 10-4 M = 5.00 (A to D); [Diol] 9 10-3 M = 5.0 (A to D); [RhCl3] 9 10-5 M = 2.39 (A, B),
1.40 (C and D); [H2SO4] 9 M = 1.00 (A, B), 0.30 M (C, D)
Rh(H2O)36þ + OHꢁ ꢀ Rh(OH)(H2O)52þ + H2O
Reactive species of rhodium(III)
Rh(OH)2ðH2O)þ4 + OHꢁ ꢀ Rh(OH)3ðH2O)3 + H2O
The rhodium(III) aqua ion [Rh(H2O)6]3? is quite stable,
Rh(OH)3ðH2O)3 + Clꢁ ꢀ RhCl(OH2ÞðH2O)3 + OHꢁ
and exchange between the inner sphere H2O protons and
bulk water is highly pH dependent [5]. It is also known that
RhCl2ðOH)(H2O)3 + Clꢁ ꢀ RhCl3ðH2O)3 + OHꢁ
this species is extensively hydrolyzed in water [6].
Between the species RhCl63-, which is obtained in excess
hydrochloric acid, and [Rh(H2O)6]3?, which is formed in
boiling aqueous solutions of rhodium(III) chloride, there
are several intermediate species, as follows;
Fritz [7] identified only one complex, [RhCl5(H2O)]2-
,
in 40% hydrochloric acid. On diluting the solution, after
24 h, [RhCl4(H2O)2]- and [RhCl3(OH)(H2O)2]- were also
observed. Capillary zone electrolysis has revealed the
123

742
Transition Met Chem (2011) 36:739–746
K1
existence of three species of rhodium, namely [RhCl4-
(H2O)2]-, [RhCl3(H2O)3] and [RhCl2(H2O)4]?, on changing
the pH of the solution from 1.0 to 3.50 with hydrochloric acid,
along with a decrease in the cationic form and increase in the
uncharged complex with the formation of two negatively
charged chlorocomplexes [8]. The present study was carried
out in acidic medium, and the positive effect of chloride ions
on the rate indicates that both [RhCl4(H2O)2]- and
[RhCl5(H2O)]2- may be present in the reaction mixture.
The absence of saturation kinetics even at high concentrations
of catalyst indicates that the catalyst is not forming a complex
with the organic substrate. The positive effect of chloride
ions on the rate suggests that [RhCl4(H2O)2]-converts into
[RhCl5(H2O)]2-, which ultimately catalyzes the reaction.
Ce4þ + HSO4ꢁ ꢀ CeSO42þ + Hþ
K2
ð1Þ
ð2Þ
ð3Þ
CeSO24þ + HSO4ꢁ ꢀ Ce(SO4Þ2 + Hþ
Ce(SO4Þ2 + HSOꢁ4 ꢀ Ce(SO4Þ3 + Hþ
K3
2ꢁ
The equilibrium constants K1, K2, and K3 for the above
steps are 3,500, 200 and 20, respectively, at 25 °C. Under our
experimental conditions, we deduce that total cerium(IV) is
mainly present as Ce(SO4)2. The concentration of Ce4?
species in a solution having [Ce(IV)] = 0.00125 M and
[H2SO4] = 1.0 M may be calculated from Eq. 4, which has
been derived from Eqs. 1 and 2. The value was found to be
1.0 9 10-9 M.
½Ce(IV)ꢂTotal
ꢀ
ꢁ
Oxidative cleavage of diols and related compounds
Â
Ã
Â
Ã
Â
Ã
2
2
¼ Ce4þ 1þK1 HSO4ꢁ =½HþꢂþK1K2 HSOꢁ4 =½Hþꢂ ð4Þ
Many oxidants readily cleave 1,2-glycols [9, 10]. Cleavage
of 1,2-glycols by lead(IV) probably involves the formation
of a bidentate metal glycol complex which then breaks down
to products via a two-electron process, which gets support
from the failure to trap radical intermediates and the fact that
cis diols and threo diols are more readily oxidized than
the trans diols and erythro diols, respectively [11]. On the
other hand, cleavage of 1,2-glycols by cerium(IV) seems to
involve coordination of only one hydroxy group followed by
a one-electron oxidation to give an intermediate radical.
This is supported by the similar rates of oxidation of 1,2-
glycols and their monomethyl ethers and also by trapping
experiments [10, 12]. Moreover, based on orbital symmetry
considerations, the former cyclic mechanism is a forbidden
process for one-electron oxidants such as cerium(IV) [13].
The cleavage of 1,2-glycols by cerium(IV) is just a special
case of the general oxidative cleavage of alcohols by
cerium(IV) [14, 15]. The same mechanism seems to operate
for all these oxidative cleavages, and as expected, the
a-hydroxy radicals are sufficiently stable that alcohols
which can form them (i.e., 1,2-glycols) undergo rapid
cleavage. The polar nature of the transition state [14] also
favors the cleavage of a-hydroxy radicals. In the uncata-
lyzed oxidation of diols, in which hydroxyl groups are not at
the vicinal position, only the corresponding hydroxy alde-
hydes were obtained as oxidation products [16]. It has been
suggested that oxidation of diols is followed by the scission
of a C–H bond instead of C–C bond [17], which has been
reported previously in the case of glycols [18].
The range of concentration of acid in which the present
study was carried out and the steep fall in rate with
increasing concentration of sulfuric acid indicates that the
concentration of other species would be negligible. Thus,
Ce(SO4)2 has been taken as the reactive species of CeIV in
aqueous sulfuric acid medium, as considered by other
workers also [17].
Mechanism
Cerium(IV) forms a complex with the organic substrate,
which then decomposes to a free radical. Transfer of the
electron of the radical to the catalyst initiates breaking of
the C–H bond, resulting in the formation of the organic
product plus rhodium(II). Rhodium(II) is quickly oxidized
back to rhodium(III) by cerium(IV).
According to the rate determining step of Scheme 1, the
rate in terms of [CeIV] is given as
d½CeIVꢂ
ꢁ
¼k½C]
ð5Þ
dt
where [C] is concentration of the complex. The total
concentration of catalyst in the reaction mixture may be
given as
½RhIIIꢂ = [RhCl4ðH2O)2ꢂꢁ + [RhCl5ðH2O)]2ꢁ + [C] ð6Þ
T
where [RhIII]T is the total concentration of the catalyst. The
concentrations of [RhCl4(H2O)2]-, [Sꢀ] and [RhCl5(H2O)]2-
calculated from Scheme 1 may be given as
Reactive species of cerium(IV) sulfate
½RhCl4ðH2O)ꢁꢂ = [RhCl5ðH2OÞ2ꢁꢂ½H2O]/K1½C]
ð7Þ
ð8Þ
ð9Þ
2
½Sꢀꢂ¼ K2½S][CeIVꢂ=½CeIIIꢂ½Hþꢂ
Cerium(IV) forms a number of complexes in sulfuric acid
solution. Hardwick and Robertson [19] have reported the
following equilibria in sulfuric acid solutions of 2 M ionic
strength;
½RHCl5ðH2O)2ꢁꢂ = [C]/K3½Sꢀꢂ
Substituting the value of [Sꢀ] from Eq. 8 into Eq. 9,
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Transition Met Chem (2011) 36:739–746
743
Scheme 1 Probable oxidation
path for 2-methyl pentane-2,4-
diol
K1
[RhCl5(H2O)]2- + H2O
O
RhCl4(H2O)21- + Cl-
OH
(I)
IV
Ce
OH
OH
K2
O
OH
+ H+
+ CeIII
(II)
+ CeIV
III
Rh
O
OH
K3
O
RhIII
OH
+
(III)
(complex C)
III
Rh
O
O
OH
OH
k
(complex C)
+ H+ + RhII
(IV)
(V)
Slow
RhII + CeIV
RhIII + CeIII
Fast
½Cꢂ½CeIIIꢂ½Hþꢂ
d½CeIVꢂ
2kK1K2K3½Sꢂ½CeIVꢂ½Clꢁꢂ½RhIIIꢂT
½RhCl5ðH2O)2ꢁ
¼
ð10Þ
ꢁ
¼
ð14Þ
K2K3½S][CeIVꢂ
½CeIIIꢂ½HþꢂþK1K2K3½S][CeIVꢂ½Clꢁꢂ
dt
Substituting the values of [RhCl4(H2O)-2 ] from Eq. 7
and [RhCl5(H2O)2-] from Eq. 9 into Eq. 6, the total
concentration of catalyst in terms of [C] may be given as
At low concentrations of oxidant and the organic substrate,
the inequality [CeIII][H?] ꢃ K1K2K3[S][CeIV][Cl-] is valid
and the rate law Eq. 14 becomes
½Cꢂ½H2Oꢂ½CeIIIꢂ½Hþꢂ
½Cꢂ½CeIIIꢂ½Hþꢂ
d½CeIVꢂ 2kK1K2K3½Sꢂ½CeIVꢂ½Clꢁꢂ½RhIIIꢂ
½Rh(III)]T¼
þ
+ [C]
T
K1K2K3½Clꢁꢂ½S][CeIVꢂ K2K3½S][CeIVꢂ
ꢁ
¼
ð15Þ
½CeIIIꢂ½Hþꢂ
This rate law explains the nature shown by all the
dt
ð11Þ
From this equation, the concentration of complex C
comes out to be
reactants at their low concentrations, i.e., first-order kinetics
with respect to oxidant, organic substrate, catalyst and the
chloride ions. The retarding effects of cerium(III) and H?
are also explained by this equation. At higher concentrations
of oxidant and organic substrate, the reverse inequality
[CeIII][H?] ꢄ K1K2K3[S][CeIV][Cl-] becomes valid, and
Eq. 14 may be written as
K1K2K3½Sꢂ½CeIVꢂ½Clꢁꢂ½RhIIIꢂT
½C] =
½CeIIIꢂ½Hþꢂf [H2O] + K1½Clꢁꢂg + K1K2K3½S][CeIVꢂ½Clꢁꢂ
ð12Þ
Now, on putting the concentration of the complex C as
obtained from Eq. 12 into Eq. 5, the rate in terms of
concentration of cerium(IV) may be given by Eq. 13. This
equation includes a factor of 2 because two moles of
oxidant are consumed to regenerate the catalyst in its
original state, thus
d½CeIVꢂ
ꢁ
¼ 2k½RhIIIꢂT
ð16Þ
dt
Equation 16 explains the observed first-order kinetics
with respect to catalyst over a very wide range of
concentrations. This equation also predicts that at higher
concentrations of H? or cerium(III), the rate becomes
independent of concentration. However, contrary to the
experimental findings, according to Eq. 16, the rate with
respect to both oxidant and chloride should also become
independent of concentration at higher concentrations.
Here, it is important to note that the effect of chloride on
the rate was studied by adding these ions externally in the
form of a solution of potassium chloride. Thus, the effect
of further increase in chloride concentration could not be
d½CeIVꢂ
ꢁ
dt
2kK1K2K3½Sꢂ½CeIVꢂ½Clꢁꢂ½RhIIIꢂT
¼
½CeIIIꢂ½Hþꢂf [H2O] + K1½ClꢁꢂþK1K2K3½S][CeIVꢂ½Clꢁꢂg
ð13Þ
As the study was carried out in aqueous medium, the
inequality [H2O] ꢃ K1[Cl-] may be considered valid and
we get
123

744
Transition Met Chem (2011) 36:739–746
studied because the reaction becomes too fast to be studied
conveniently. It may be possible that had the study been
conducted at still higher chloride concentrations, the rate
would have become independent of [Cl-]. Further, the
effect of the oxidant concentration is somewhat surprising
because at higher concentrations instead of showing zero-
order kinetics, the rate starts decreasing. It is interesting
to note that the decrease in rate starts only when the
concentration of oxidant exceeds 6.0 9 10-4 M in all the
cases. It is possible that the decrease may be due to
formation of some dimeric species of Ce(IV) in the
reaction mixture. Formation of dimeric species of Ce(IV)
has been reported by a number of workers in sulfuric, nitric
[16, 20] and perchloric acids [21], while the existence of a
trimeric cerium species has been reported in acetic acid
medium [22]. In a number of studies, it has been reported
that the first-order rate constant decreases with increasing
cerium(IV) concentrations due to the formation of dimeric
species [23–25]. The retarding trend shown by the oxidant
at higher concentrations may also be due to the formation
of a complex between the oxidant and sulfuric acid,
according to the following equilibria;
and 3-methyl pentane-2,4-diol, respectively. Since the
rates were measured at the beginning of the reaction, the
concentration of the product cerium(III) at the point where
the rates were measured will be negligible. Similarly, in the
absence of externally added chloride, the concentration
of chloride will be only due to the traces of catalyst present
in the reaction mixture (in the range of 10-5 M) and this
may also be ignored while calculating the rate constants.
The kinetic measurements were carried out at four
temperatures. The values of activation energies 18.30
( 0.03), 14.82 ( 0.05), 11.44 ( 0.44) and 20.52 ( 0.01)
(kJ mol-1), calculated with the help of the Arrhenius plots,
were used to calculate entropy of activation and free energy
of activation values which were found to be -44.08,
-44.51, -45.03 and -43.86 (JK-1 mol-1); and 31.74,
28.39, 25.13 and 33.90 for butane-1,3-diol, butane-1,4-diol,
2-methyl pentane-2,4-diol and 3-methyl pentane-2,4-diol,
respectively.
Experimental
Cerium(IV) sulfate, diols (CDH), sulfuric acid, acetic acid,
ferroin (Merck) and cerium(III) sulfate (Fluka) were used
as supplied by preparing their solutions in double distilled
water. The concentration of rhodium(III) chloride (Johnson
Matthay) was 4.78 9 10-3 M, which was prepared by
dissolving the sample in a minimum amount of AR HCl
(5.0 9 10-2 M). Ferrous ammonium sulfate solution was
standardized against a standard solution of potassium
dichromate (Merck). Cerium(IV) sulfate was prepared by
dissolving the sample in 17.0 N sulfuric acid and was
standardized with standard ferrous ammonium sulfate. All
other chemicals used were either AR or chemically pure
substances. The progress of the reaction was measured
(constant temperature 0.1 °C) at different time intervals
by transferring aliquots to a fixed amount of ferrous
ammonium sulfate solution (in slight excess to cerium (IV)
sulfate initially taken) and estimating the remaining
ferrous ammonium sulfate, with standard cerium(IV) sul-
fate solution using ferroin as indicator. In this way, the
remaining amount of ferrous ammonium sulfate directly
corresponds to the amount of cerium(IV) consumed in the
reaction mixture. In all kinetic runs, diol concentration was
kept in excess.
K0
Ce(SO4Þ2 + HSOꢁ4 ꢀ HCe(SO4Þ3ꢁ
ð17Þ
K00
HCe(SO4Þꢁ + HSO4ꢁ ꢀ H2Ce(SO4Þ
Ce(SO4Þ2 = [H2Ce(SO4Þ4ꢂ =½HSOꢁ4 ꢂ
[H2Ce(SO4)4]2-, a well-known unreactive species of
cerium(IV) sulfate in sulfuric acid medium [26, 27], would
retard the rate as given in Eq. 19. Similar behavior has
been reported in the oxidation of hydrocarbons by
cerium(IV) [28].
ð18Þ
ð19Þ
2ꢁ
3
4
2ꢁ
2
Further verification of the final rate law Eq. 14 may be
given by rewriting the equation as
1
½CeIIIꢂ½Hþꢂ
1
¼
þ
ꢁd[CeIVꢂ=dt 2kK1K2K3½S][CeIVꢂ½Clꢁꢂ½RhIIIꢂT 2k½RhIIIꢂT
ð20Þ
From this equation, values of k calculated from the
intercepts of the graphs of 1/rate versus 1/[S] and [H?]
were obtained as 38.02 ( 0.04) and 20.90 ( 0.00); 90.90
( 0.02) and 93.45 ( 0.09); 27.77 ( 0.12) and 24.63 ( 0.
10); 25.51 ( 0.07) and 10.20 ( 0.00) for butane-1,3-diol,
butane-1,4-diol, 2-methyl pentane-2,4-diol and 3-methyl
pentane-2,4-diol, respectively. Similarly, the values of
kK1K2K3 from the slopes of the graphs between 1/rate
versus 1/[S] and [H?] were calculated as 6.27 ( 0.07) and
6.15 ( 0.06); 4.74 ( 0.31) and 5.03 ( 0.00); 4.58 ( 0.10)
and 4.16 ( 0.03); 5.81 ( 0.02) and 5.71 ( 0.11) for
butane-1,3-diol, butane-1,4-diol, 2-methyl pentane-2,4-diol
Determination of kinetics
In all the kinetic runs, [diol] ꢃ [oxidant] as per Ostwald’s
isolation method, and thus, the calculated orders are given
with respect to oxidant concentrations. In all the cases,
the rate values (-dc/dt) were calculated at a fixed initial
concentration of oxidant except in the case of oxidant
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Transition Met Chem (2011) 36:739–746
745
MeCH(OH)CH(R)CR(OH)Me + 2CeIV
! MeCOCHRCR(OH)Me + 2CeIII + 2Hþ
variation, where the rate values were calculated at a fixed
initial time from the individual plots. Rate values (-dc/dt),
obtained from the initial slopes of individual plots of the
residual concentrations of cerium(IV) at various time
intervals, were finally plotted against the concentration of
the particular reactant for which the order of the reaction
was to be obtained. Orders, with respect to various reac-
tants, were confirmed by plotting log (a–x) versus time
(oxidant variation), by plotting -dc/dt values versus con-
centration of the reactant, by calculating the slope of the
double logarithmic graphs of rate versus concentration and
by calculating the rate constant for concentration (catalyst
variation). The effects of change of the concentration of
Ce(III) and Cl- on the reaction velocity were studied by
external addition of these ions. The tables and figures give
the initial concentrations of the reactants. Studies could not
be made at constant ionic strength, due to large volumes of
KCl required to keep the ionic strength constant. However,
the effect of change of ionic strength on the rate was
studied separately using a standard solution of KCl.
ð21Þ
Conclusion
In the rhodium(III)-catalyzed oxidation of non-vicinal
diols, scission of a C–H bond instead of a C–C bond takes
place, as in the case of their uncatalyzed oxidations. The
corresponding hydroxy ketone is found to be the oxidation
product. This is not the case in vicinal diols where the C–C
bond is broken [33]. The present study also shows that
rhodium(III) chloride which is a sluggish catalyst in alka-
line medium can act as an efficient catalyst in acidic
medium.
References
1. Tandon PK, Sahgal S, Singh AK, Kumar S, Dhusia M (2006) J
Mol Catal A Chem 258:320
2. Tandon PK, Dhusia M (2009) Trans Metal Chem 34:121
3. Tandon PK, Dhusia M, Sahgal S (2009) J Mol Catal A Chem
301:146
Product study and stoichiometry
Separation of the products to check whether the system can
be used for synthetic purposes and stoichiometry of the
reaction were studied by taking cerium(IV) sulfate in
large excess compared to the organic substrate in different
ratios. These reaction mixtures were kept overnight at
room temperature. After completion of the reaction, the
solution was extracted with diethyl ether (3 9 25 mL).
After evaporating the solvent, the hydrazone of the organic
product was prepared by standard methods [29]. Yellow-
ish-red precipitates of the corresponding 2,4-dinitrophenyl
hydrazones of 3-hydroxy butanal, 4-hydroxy butanal, 4-
hydroxy-4-methyl pentan-2-one and 4-hydroxy-3-methyl
pentan-2-one were dried and weighed, giving 22 and 18%
yields of corresponding hydroxy ketones. The final oxida-
tion products were confirmed by spot tests [30], chroma-
tography [31] and spectroscopy (Supplementary Material 1
4. Tandon PK, Gayatri SahgalS, Srivastava M, Singh SB (2007)
Appl Organomet Chem 21:135
5. Lee JD (1996) Concise Inorganic Chemistry, 5th edn. Chapman
and Hall, London, p 797
6. Cotton FA, Wilkinson G, Murillo CA, Bochmann M (1999)
Advanced Inorganic Chemistry, 6th edn. Wiley-Interscience,
New York, p 1048
7. Thornton MJ, Fritz JS (1997) J Chromatogr A 770:301
8. Aleksenko SS, Gumenyuk AP, Mushtakova SP, Kozhina LF,
Timerbaev AR (2003) Talenta 61:195
9. Perin AS (1969) In: Augustine RL (ed), Oxidation, Marcel
Dekker, New York, NY, Vol I, Chap IV
10. Trahanovsky S, Young LH, Bierman MH (1969) J Org Chem
34:869
11. Bunton CA (1965) In: Wiberg KB (ed) Oxidation in Organic
Chemistry, Part A; Academic Press, New York, NY, Chap VI
12. Littler JS, Waters WA (1960) J Chem Soc 2767
13. Littler JS (1971) Tetrahedron 27:81
14. Nave PM, Trahanovsky WS (1971) J Am Chem Soc 93:4536
15. Trahanovsky WS, Cramer JJ (1971) Org Chem 36:1890
16. Prakash A, Mehrotra RN, Kapoor RC (1979) J Chem Soc (Dalton
Trans) 205
17. Sankhla PS, Mehrotra RN (1973) J Inorg Nucl Chem 35:891
18. Richardson WH (1965) In: Wiberg KB (ed) Oxidation in Organic
Chemistry, Part A; Academic Press, New York, NY, p 244
19. Hardwick TJ, Robertson E (1951) Can J Chem 29:828
20. Ardon M, Stein G (1956) J Chem Soc 104
21. Dorfman MK, Gryder JW (1962) Inorg Chem 1:799
22. Ford PC, Wiberg KB (1968) Inorg Chem 7:369
23. Sankhla PS, Mehrotra RN (1976) Indian J Chem 14A:663
24. Rao GG, Rao BM (1972) Anal Chim Acta 59:461
25. Sankhla PS, Mehrotra RN (1975) Indian J Chem 13:904
26. Tandon PK, Krishna B (1985) Kinetika i Kataliz 26:607
27. Singh MP, Singh HS, Verma MK (1980) J Phy Chem 84:256
28. Tandon PK, Srivastava M, Kumar S, Singh S (2009) J Mol Catal
A Chem 304:101
to 4). A broad CO–H str was observed at 3,550–3,200 cm-1
,
indicating associated O–H due to hydrogen bonding, while
hydrogen bonded carboxyl peaks (tC=O str. peaks at 1,631,
1,634, 1,632 and 1,631 Cm-1 [32]. The tC–Ostr. bands
were at 1,197, 1,193, 1,200 and 1,202 Cm-1, while t-CH3str
.
(C–H asymmetric and symmetric) were observed at 1,343,
1,350; 1,348, 1,382; 1,348, 1,387 and 1,343, 1,383 for
butane-1,3-diol, butane-1,4-diol, 2-methyl pentane-2,4-diol
and 3-methyl pentane-2,4-diol, respectively. The stoichi-
ometry of the reaction as determined by estimating the
unconsumed cerium(IV) in the reaction mixture after
complete oxidation of the calculated quantity of the organic
substrate may be given by Eq. 21 where R = H or Me;
123

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Transition Met Chem (2011) 36:739–746
29. Shriner RL, Hermann CKF, Morrill TC, Curtin DY, Fuson RC
(1998) The systematic identification of organic compounds.
Wiley, New York, p 321
30. Feigl F (1960) Spot tests in organic chemistry. Elsevier, New
York, p 369
31. Mikes O (1966) Laboratory handbook of chromatographic
32. Silverstein RM, Bassler GC, Morill TC (1981) Spectrometric
identification of organic compounds, 4th edn. Wiley, New York,
p 120
33. Trahanovsky WS, Gilmore JR, Heaton PC (1973) J Org Chem
38:1973
methods, D. Van Nostrand Co. Ltd., London (Eng. Ed.) Chap 2,
p 96
123