Kinetics of the oxidation of heterocyclic aldehydes by quinolinium dichromate

Article abstract of DOI:10.1246/bcsj.75.2215

Quinolinium dichromate in sulfuric acid oxidizes heterocyclic aldehydes (2-furaldehyde, 2-pyrrolecarbaldehyde, 2-thiophenecarbaldehyde) to the corresponding acids in a 50% (v/v) acetic acid-water medium. The kinetic data on the rates of oxidation of the substrates have been discussed with reference to the aldehyde hydration equilibria. The kinetic results support a mechanistic pathway proceeding via a rate-determining oxidative decomposition of the chromate ester of the aldehyde hydrate.

Full text of DOI:10.1246/bcsj.75.2215

c. Jpn.
© 2002 The Chemical Society of Japan
Bull. Chem. Soc. Jpn., 75, 2215–2220 (2002) 2215
e Chemical
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Kinetics of the Oxidation of Heterocyclic Aldehydes by Quinolinium
Dichromate
Oxidation of Heterocyclic Aldehydes
G. S. Chaubey et al.
Girija Shankar Chaubey, Simi Das, and Mahendra Kumar Mahanti*
02
15
20
SJA8
09-2673
060
Department of Chemistry, North-Eastern Hill University, Shillong 793022, India
(Received March 8, 2002)
Quinolinium dichromate in sulfuric acid oxidizes heterocyclic aldehydes (2-furaldehyde, 2-pyrrolecarbaldehyde, 2-
thiophenecarbaldehyde) to the corresponding acids in a 50% (v/v) acetic acid–water medium. The kinetic data on the
rates of oxidation of the substrates have been discussed with reference to the aldehyde hydration equilibria. The kinetic
results support a mechanistic pathway proceeding via a rate-determining oxidative decomposition of the chromate ester
of the aldehyde hydrate.
02
02
.1
3.6
730 cm−1, characteristic of the dichromate ion. Acetic acid (SD,
AR grade) was distilled before use. Sulfuric acid (E. Merck) was
used after a check of its physical constants. IR spectra were re-
corded on a FT-IR (DA- 8, Bomen) spectrophotometer.
The oxidation of heterocyclic aldehydes by Mn(Ⅶ) had
highlighted the kinetic aspects and nature of the formed prod-
uct.1 Cerium(Ⅳ) ion in an acetic acid (25%, v/v) medium was
used to oxidize 2-furaldehyde to the acid; the reaction involved
a free-radical mechanism.2 Thallium(Ⅲ) was used for the oxi-
dation of 2-furaldehyde in a perchloric acid solution, and it
was observed that the rate of the reaction did not depend on the
concentration of H+ ions.3 The kinetics of the oxidation of
heterocyclic aldehydes by bromic acid in a H2SO4–HOAc me-
dium showed that the reaction was first-order both in the oxi-
dant and substrate concentrations, but showed a second-order
dependence on the concentration of H+ ions.4 Kinetic studies
on the oxidation of 2-furaldehyde by chromic acid5 and quino-
linium chlorochromate6 in acetic acid water media showed a
first-order dependence each on the concentrations of the sub-
strate, oxidant and hydrogen ions.
In the oxidation of heterocyclic aldehydes, there exists the
possibility of the reaction taking place either at the heteroatom
or at the aldehydic function. The aims of the present investiga-
tion were: (a) to highlight the effect of the heteroatom on the
rate of the reaction and (b) to determine the site of attack of the
oxidant. For this purpose, we carried out a kinetic study of the
oxidation of heterocyclic aldehydes (2-furaldehyde, 2-pyrrole-
carbaldehyde and 2-thiophenecarbaldehyde) by a chromi-
um(Ⅵ) reagent, quinolinium dichromate [QDC, (C9H7N+H)2-
Cr2O72−], in an acid medium, in 50% acetic acid–water (v/v),
under a nitrogen atmosphere. This study forms part of our
continuing efforts concerning the oxidation of organic sub-
strates by quinolinium dichromate in general,7 and aldehydes
in particular.8
The method used for the kinetic determinations was described
earlier.7
Product Analysis. Thirty cm3 of water was taken and cooled
in ice. Concentrated H2SO4 (7.9 g, 0.08 mol dm−3) was added
slowly with constant cooling. When the acid solution had cooled
to room temperature, quinolinium dichromate (QDC 9.52 g, 0.02
mol dm−3) was added and the mixture was warmed to 313 K for
complete dissolution of the QDC. To this mixture, 0.015 mol dm−3
of a substrate (1.45 g of 2-furaldehyde, 1.43 g of 2-pyrrolecarbal-
dehyde, and 1.69 g of 2-thiophenecarbaldehyde), taken in 25 cm3
of 50% acetic acid-water solution, was added. The reaction mix-
ture was stirred at 313 K for 48 h under nitrogen. The organic lay-
er was extracted thrice with ether (25 cm3 each time), and the
combined organic extracts were washed with water and dried over
anhydrous Na2SO4. The oxidized products (2-furancarboxylic
acid from 2-furaldehyde; 2-pyrrolecarboxylic acid from 2-pyrrole-
carbaldehyde; and 2-thiophenecarboxylic acid from 2-thiophene-
carbaldehyde) were obtained after the complete removal of ether
(melting points were in agreement with literature values ; yields ꢀ
85–90%). Each reaction product was subjected to IR (KBr) analy-
sis, and characterized as follows: (i) 2-Furancarboxylic acid:ν =
3000, 2860 (br, s, –OH), 2583, 1690 (s, CwO), 1470, 1305, 1245,
1020, 930, 760 cm−1, (ii) 2-Pyrrolecarboxylic acid:ν = 3000,
2850 (br, s, –OH), 1692 (s, CwO), 1545, 1355, 1105, 925, 750 cm−1
,
(iii) 2-Thiophenecarboxylic acid: ν = 3090, 2850 (br, s, –OH),
2621, 1690 (s, CwO), 1530, 1350, 1100, 910, 750 cm−1
.
Results and Discussion
Experimental
The oxidation of heterocyclic aldehydes (2-furaldehyde, 2-
pyrrolecarbaldehyde, 2-thiophenecarbaldehyde) by QDC re-
sulted in the formation of the corresponding acids. Under the
present experimental conditions, there was no further oxida-
tion of the acids.
Materials and Methods. 2-Furaldehyde (S.D. fine chemicals
co.) and 2-thiophenecarbaldehyde (Aldrich) were purified by dis-
tillation under reduced pressure. 2-Pyrrolecarbaldehyde (Aldrich)
was recrystallized before use. The oxidant, quinolinium dichro-
mate [QDC, (C9H7N+H)2Cr2O72−], was prepared by the reported
method,9 and its purity was checked by spectral analysis. The in-
frared spectrum (KBr) exhibited bands at 930, 875, 765 and
The stoichiometry of the reaction was determined.7 Stoichi-
ometric ratios, ∆[QDC]/∆[substrate], in the range 0.66 to 0.69
were obtained, which conformed to the following overall

2216 Bull. Chem. Soc. Jpn., 75, No. 10 (2002)
Oxidation of Heterocyclic Aldehydes
equation:
kalinity.10 Hence, it would be justified to propose that in the
range of acid concentrations used, the oxidant QDC was con-
verted to the protonated chromium(Ⅵ) species. Earlier reports
have established the involvement of protonated Cr(Ⅵ) species
in chromic acid oxidation reactions.11,12
3 C5H4O2 + 2 Cr + 3 H2O → 3 C5H4O3 + 2 Cr + 6 H+ (1)
Ⅵ
Ⅲ
(2-furaldehyde).
Using pseudo-first-order conditions, individual kinetic runs
were observed to be first order in QDC. The pseudo-first-order
rate constants (k) were independent of the initial concentration
of the oxidant (Table 1). The order of the reaction with respect
to the substrate concentration was obtained by varying the al-
dehyde concentration and observing the effect on the rate at
constant [QDC] and [H+]. The results are given in Table 1.
The order with respect to the concentration of acid was ob-
tained at constant aldehyde and QDC concentrations. The data
given in Table 1 indicate a first-order dependence on the con-
centration of the acid. Under the acid concentrations used in
the present investigation, the protonation of the aldehydes
would be less significant. The possibility of the aldehydes get-
ting protonated can be ruled out on the basis that aldehydes are
extensively hydrated in an aqueous medium, and are present as
equilibrium mixtures of the carbonyl and hydrated forms. The
formation constants are thus not dependent on the acidity or al-
The oxidation of the substrates by QDC was studied over
the temperature range 303–323 K; the rate data and activation
parameters are given in Table 2. The negative values of ∆S*
provided support for the formation of a rigid activated com-
plex.
The oxidation of the substrates was studied in solutions con-
taining varying proportions of water and acetic acid. The di-
electric constants (D) of water–acetic acid mixtures were cal-
culated from the dielectric constants of the pure solvents (at
313 K: water = 73.28, acetic acid = 6.29).13 It was observed
that there was an increase in the rate of the reaction with a de-
crease in the dielectric constant of the medium (Table 3).
There was no induced polymerization of acrylonitrile or a
reduction of mercury(Ⅱ) chloride,14 which indicated that a one-
electron oxidation was quite unlikely. Control experiments,
performed in the absence of a substrate, did not show any ap-
preciable change in the concentration of QDC.
Table 1. Rate Data for the Oxidation of Heterocyclic Aldehydes at 313 K
102 [Substrate]
mol dm−3
103 [QDC]
mol dm−3
[H2SO4]
mol dm−3
104 k/s−1
2-Furalde-
hyde
1.25
2-Pyrrolecarb-
aldehyde
1.18
2-Thiophenecarb-
aldehyde
1.14
1.0
2.5
5.0
7.5
10
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.75
1.0
1.25
1.50
3.12
6.21
9.32
12.5
1.22
1.25
1.24
1.27
1.88
2.50
3.20
3.80
2.91
5.90
8.71
11.8
1.15
1.18
1.19
1.14
1.75
2.36
2.90
3.55
2.83
5.67
8.50
11.2
1.14
1.13
1.15
1.12
1.64
2.20
2.70
3.40
1.0
0.75
0.50
0.25
0.10
1.0
1.0
1.0
1.0
Table 2.
Aldehydesb) by QDC
Temperature and Activation Parametersa) for the Oxidation of Heterocyclic
104 k/s−1
Tempc)/K
2-Furaldehyde 2-Pyrrolecarbaldehyde 2-Thiophenecarbaldehyde
303
308
313
318
323
0.62
0.94
1.25
1.78
2.54
0.59
0.89
1.18
1.74
2.41
0.56
0.84
1.14
1.71
2.35
∆H* (kJ mol−1
)
51
−152
54
−145
56
−142
∆S* (J mol−1 K−1
)
a) Error limits:∆H* 2 kJ mol−1; ∆S* 5 J mol−1 K−1
.
b) [Substrate] = 1.0 × 10−2 mol dm−3; [QDC] = 1.0 × 10−3 mol dm−3; [H2SO4] =
0.5 mol dm−3
c) 0.1 K.
.

G. S. Chaubey et al.
Bull. Chem. Soc. Jpn., 75, No. 10 (2002) 2217
Table 3. Solvent Effect for Oxidation of Heterocyclic Aldehydesa) by QDC at 313 K
Dielectric
constants
D
104 k/s−1
H2O:AcOH
(%, v/v)
2-Furaldehyde
2-Pyrrolecarbaldehyde
2-Thiophenecarbaldehyde
50:50
45:55
40:60
35:65
30:70
39.79
36.44
33.09
29.74
26.39
1.25
1.82
2.51
3.80
6.30
1.18
1.58
2.09
2.95
4.27
1.14
1.41
1.78
2.51
3.55
a) [Substrate] = 1.0 × 10−2 mol dm−3; [QDC] = 1.0 × 10−3 mol dm−3; [H2SO4] = 0.5 mol dm−3
.
Mechanism. The kinetic results showed that the rate of
oxidation of heterocyclic aldehydes was dependent on the first
powers of the concentrations of each (substrate, oxidant and
acid). The acid catalysis of the reaction must be related to the
structure of the oxidant (QDC), which was converted to a pro-
tonated species at the concentrations of the mineral acid used.
Quinolinium dichromate (QDC) is a dimetallic species, an an-
ionic condensed form of chromic acid. Aqueous solutions of
H2O, a decrease in the dielectric constant of medium (increase
in the acetic acid content) would favor the dichromate form
over the chromate form. If ion-pairs were to be formed in this
medium, it would be expected that they would have a higher
ion-pair association constant for the dichromate ion, which
would again favor the dichromate ion. The absence of any salt
effects on the rate of oxidation indicated that the reaction was
not of the ion-ion type. If the reaction was assumed to involve
two neutral molecules, then the plot of log k versus (D−1)/(2D
+ 1) would have been linear; this was not found to be so. Al-
though the range of dielectric constants used for these reac-
tions was not large, plots of log k versus 1/D were found to be
linear, with positive slopes, which indicated that the reactions
were of the ion-dipole type.18
It has been shown that aldehydes are extensively hydrated in
aqueous solutions and many oxidation reactions proceed via
the hydrate form.10,19–24 Table 4 lists the experimental rate
constants (k) for the oxidation of the aldehydes by QDC. The
aldehyde hydrate dissociation constants (Kd) pertaining to the
reaction
−
chromic acid contain ions such as CrO42−, HCrO4 and
Cr2O72−, besides other protonated species such as H2Cr2O7,
−
HCr2O7 and H2CrO4.15 The ionization constant for the
HCrO4− ion, HCrO4− ꢀ H+ + CrO42−, was 3.0 × 10−7 mol/L;
hence, in dilute aqueous acid, the concentration of CrO42− ions
was negligible. This has been amply substantiated by Michel
et al.,16 who examined the Raman spectra of chromate, dichro-
mate and chlorochromate species, and found that the protonat-
−
ed form of chromate HCrO4 did not exist in aqueous solu-
tions of Cr(Ⅵ) compounds. The ionization constant for the
−
−
2−
HCr2O7 ion given by HCr2O7 ꢀ H+ + Cr2O7 was 0.85
mol/L; hence, in solutions where pH ꢀ 1, the ionization may
be considered essentially complete. Consequently, of all the
ions involving hexavalent chromium, the only ones present in
large concentrations in solutions of dilute mineral acid would
Kd
(2)
RCH(OH)2
RCHO + H2O
be HCrO4 and Cr2O72−. These ions are in equilibrium with
−
are also given. From k and Kd, two sets of rate constants for
the oxidation of the aldehyde in only one of the forms present
in solution were computed. The values of kHy were obtained
by assuming that only the hydrate form appears in the rate law,
−
each other according to the following equation: 2 HCrO4
ꢀ
2−
Cr2O7 + H2O, with the value of Kd = 35.5. According to
this equilibrium, an increase of the hydrochromate concentra-
tion should be significant with dilution. When the Raman lines
were examined under dilution, it was established that at pH =
ν = kHy[QDC][RCH(OH)2].
(3)
2−
11, the Cr(Ⅵ) ion was 100% present in the form of the CrO4
2−
ion, whereas at pH = 1.2, it was 100% as the Cr2O7 ion.16
Similarly, the values of kA were calculated using the concentra-
tion of free aldehydes according to the rate law,
Hence, at concentrations of acid larger than 0.05 M, the
dichromate ion (and its protonated forms) would be the pre-
dominant species. In aqueous solutions of K2Cr2O7, spectral
ν = kA[QDC][RCHO].
(4)
2−
studies have shown that Cr2O7 was the predominant spe-
The values of kHy and kA are given in Table 4. Using the σ+
values derived from a consideration of the electrophilic substi-
tution for the hetero systems,25 a plot of log kHy against σ+
was linear (r = 0.9632), with a slope of ρ = +3.0. On the oth-
er hand, the correlation of σ+ with kA gave a value of ρ =
+0.675 (r = 0.994). The correlation with kHy supported the
mechanistic pathway for the oxidation reactions as proceeding
through the hydrated form of the aldehydes. Hence, a mecha-
nism involving a direct hydrogen-transfer reaction between a
free aldehyde and QDC was very unlikely. Thus, the rate-ac-
celerating effect of the electronegative substituents could be
interpreted in terms of greatly increased hydration (Table 1). It
cies.17 In the present investigation, since the concentrations of
acid used were in the range 0.5 to 1.5 M (1M = 1 mol dm−3),
the dichromate ion (or its protonated form) would be the pre-
dominant species. Moreover, the protonated Cr(Ⅵ) species
would be a more reactive electrophile capable of increasing its
rate of coordination to the substrate.
The effect of a change in the solvent composition (water–
acetic acid, %, v/v) on the rate of oxidation for the substrates
was studied. It was observed that an increase in the water con-
tent of the medium showed a decrease in the rate of oxidation
(Table 3). The magnitude of this effect could be analyzed by
−
2−
suggesting that, for the equilibrium 2HCrO4 ꢀ Cr2O7
+

2218 Bull. Chem. Soc. Jpn., 75, No. 10 (2002)
Table 4. QDC Oxidation of Heterocyclic Aldehydes at 313 K
Oxidation of Heterocyclic Aldehydes
Aldehydes
2-Furaldehyde
2-Pyrrolecarbaldehyde
Kd
104 k/s−1
kHy/M−2 s−1
kA/M−2 s−1
1.29
0.92
0.84
1.25
1.18
1.14
16.56 1.28
10.99 0.41
9.53 0.17
12.84 1.28
11.94 0.45
11.34 0.20
2-Thiophenecarbaldehyde
would be pertinent to compare the ρ value of +0.675, obtained
in this study, with the ρ values obtained for a series of aromatic
aldehydes in 91% acetic acid.26,27 Since aromatic aldehydes
were hydrated only to a small extent, the similarity between
these values (ρ = +1.02 and +0.77)26,27 and our value (ρ =
+0.675) was satisfactory in that for both cases, small positive
values were obtained. The positive value of ρ could be inter-
preted as being due a superimposed effect of the ring substitu-
ents on the hydration equilibrium, wherein it had been shown
that all of the aldehydes oxidized by chromic acid were com-
pletely hydrated in aqueous solution.28 Earlier work on the ox-
idation of benzaldehyde by chromic acid had shown that a
Hammett plot of the data yielded a good straight line with a
slope ρ = +1.06 (Ref. 29). Later work on the chromic acid
oxidation of benzaldehyde had shown that the reaction pro-
ceeded by way of the chromic acid ester of hydrated benzalde-
hyde as an intermediate.30 This similarity provided additional
support for the mechanistic pathway suggested in the present
investigation, that the rate-determining step was the oxidative
decomposition of the chromate ester of an aldehyde hydrate.
The hydrated forms of the substrates would remain as undisso-
ciated molecules (since [H+] would be much greater than the
dissociation constants of the substrates), in the range of the
acid concentrations used in the present study (0.5–1.50 M). It
was therefore suggested that the protonated QDC reacted with
the hydrated form of the heterocyclic aldehydes to give the
corresponding acids.
Since aldehyde hydrates very closely resembled alcohols,
both in structure and in many aspects of oxidation, there would
be a similarity in the two oxidation reactions. The oxidation
mechanism of alcohols by chromic acid had shown that the
rate-determining step involved the decomposition of the proto-
nated acid chromate ester.31 By analogy, the oxidation of alde-
hydes would proceed via the formation of a similar intermedi-
ate (an ester of the aldehyde hydrate), which would undergo
decomposition in the rate-determining step. The hydration of
the aldehydes was not rate-determining for the oxidation reac-
tion.32 The ester of the aldehyde hydrate would be in equilibri-
um with the free aldehyde and the aldehyde hydrate, and could
be formed either by a carbonyl addition reaction to the free al-
dehyde or by esterification of the hydrate.33 It may be added
that the esterification reaction has more utility, since it helps to
understand and predict aldehyde oxidation reactions and their
relationship to the closely related oxidation of alcohols.
N and S atoms (electronegativities were: O = 3.50; N = 3.07;
S = 2.44).34 The inference was that electronegative substitu-
ents increased the oxidation rates by increasing the equilibrium
concentrations of the chromate ester of the aldehyde hydrate.
This would thus account for the order of the observed reactivi-
ty (Table 1).
The mechanistic pathway involved the formation of the es-
ter (step 1), followed by the slow oxidative decomposition of
the ester of the aldehyde hydrate (step 2). A cyclic structure
for the reaction intermediate would explain all of the features
of the oxidation reaction. The large negative entropies of acti-
vation (∆S≠) would be consistent with the formation of a cyclic
intermediate in a bimolecular reaction. If the chromium was
coordinated through the –OH group (of the aldehyde hydrate),
then the formation of the chromate ester would be facilitated.
This would increase the ease of oxidation, and conversion to
the corresponding carboxylic acid could be rationalized.
The manner of electron transfer must be established. Elec-
tron flow in a cyclic transition state has been considered,35 and
could be rationalized by assuming that if the chromium was
coordinated through the –OH group (of the aldehyde hydrate),
then the electron-transfer process could take place through the
carbon–oxygen–chromium bond. This would facilitate the for-
mation of the chromate ester, and also enhance the ease of con-
version to the product. The proton was removed in the cyclic
transition state (co-planarity of all the atoms involved), the
center of which resided on an electron-dense oxygen in the
chromate ester.36 Such a transition state envisaged the transfer
of electrons towards the chromium, occurring by the formation
of the carbon–hydrogen–oxygen bonds, as well as by the car-
bon–oxygen–chromium bonds.
The slow step of the reaction involved the participation of
the aldehyde hydrate, protonated QDC, and two electrons in a
cyclic system. Removal of the hydrogen (on the carbon) was
part of this step, as evidenced by the experimental observation
that a kinetic isotope effect was observed for the oxidation of
2-furaldehyde-d1 (kH/kD = 5.8). Since the five-membered het-
erocyclic ring system is a planar pentagon with sp2 hybridized
carbon atoms, and possesses a considerable aromatic character
arising from delocalization of the two paired electrons, it
would undergo a reaction via an electrocyclic mechanism in-
volving six electrons; being a Hückel-type system (4n + 2),
this would be an allowed process.37 This mechanism was sup-
ported by the observation that the oxidation reactions were
acid-catalyzed. Protonation of the oxidant would make it more
amenable towards nucleophilic attacks by the substrate on the
electron-deficient chromium of the oxidant.
The heteroatoms are strong resonance donors in these five-
membered ring systems, an effect which completely overrides
their inductive withdrawal. By treating these rings as transmit-
ting systems, one could look for a correlation between the
structure and the reactivity of these heterocyclic aldehydes.
The observed order of reactivity was 2-furaldehyde > 2-pyr-
rolecarbaldehyde > 2-thiophenecarbaldehyde (Table 1), which
was in conformity with the decreasing electronegativities of O,
The sequence of reactions for the oxidation of heterocyclic
aldehydes by QDC, in an acid medium is shown in the
Scheme. In an acid medium, the oxidant QDC was converted
to the protonated dimetallic chromium(Ⅵ) species (PQ) [in the
acid range used for the present investigation, the protonated

G. S. Chaubey et al.
Bull. Chem. Soc. Jpn., 75, No. 10 (2002) 2219
Scheme 1.
2−
QDC would have the Cr(Ⅵ) existing mainly as Cr2O7 (Ref.
17)]. The substrate (A) was converted to the hydrated form
(Hy). The reaction of the hydrated form (Hy) with the proto-
nated QDC (PQ) resulted in the formation of the monochro-
mate ester (E) and a Cr(Ⅵ) monomer. The monochromate es-
ter (E) underwent decomposition in the rate-determining step
to give the product (the corresponding acid), along with the
Cr(Ⅳ) species. The conversion of Cr(Ⅳ) to Cr(Ⅲ) proceeded
by a disproportionation reaction. It was shown that for the re-
action Cr(Ⅳ) + Cr(Ⅵ) → 2 Cr(ꢁ), the standard potential for
the Cr(Ⅵ)–Cr(ꢁ) couple was extremely favourable (E0 = 0.62
V),38 and this reaction would proceed rapidly. The Cr(ꢁ)–
Cr(Ⅲ) couple has a potential of 1.75 volt, which would enable
the rapid conversion of Cr(ꢁ) to Cr(Ⅲ), after the reaction of
Cr(ꢁ) with the substrate.38,39
Substituting the values of [PQ] and [Hy] in Eq. 5 (taking the
activity of water to be unity), we obtain
−d [QDC]/dt = K1K2 k3 [A][QDC][H+]
(6)
From this rate expression (Eq. 6), it is clear that the reaction
exhibited a first-order dependence with respect to the concen-
trations of each (substrate, oxidant, and acid). Hence, −2.303
d log [QDC]/dt = k = K1K2 k3 [A][QDC][H+]. This rate law
explains all of the experimentally observed results.
The data collected demonstrated that the application of
QDC to the oxidation of heterocyclic aldehydes led to the for-
mation of carboxylic acids, substantiating the mechanism of
the oxidation reaction wherein there was an attack of the oxi-
dant on the aldehydic function, leaving the heteroatom site in-
tact. While highlighting the importance of QDC as an oxidant,
this study emphasized the efficiency of the reactions of QDC
with heterocyclic aldehydes, which could prove to be a regi-
oselective route for the synthesis of carboxylic acids.
If the mechanism shown in the Scheme is correct, then the
attack of the protonated QDC (PQ) on the aldehyde hydrate
(Hy) would be crucial, and would be favored by the formation
of the cyclic chromate ester (E). Based on the mechanism
shown in the Scheme, the rate law has been derived as follows:
Financial support from the University Grants Commission,
New Delhi, under the Special Assistance Program, is gratefully
acknowledged.
−d [QDC]/dt = k3 [E] = k3 [Hy][PQ],
(5)
where [PQ] = K1 [QDC][H+], and [Hy] = K2 [A] [H2O].

2220 Bull. Chem. Soc. Jpn., 75, No. 10 (2002)
Oxidation of Heterocyclic Aldehydes
References
(1986).
17 M. Creslak- Golonka, Coord. Chem. Rev., 109, 223 (1991).
18 E. S. Amis, “Solvent Effects on Reaction Rates and
Mechanisms,” Academic Press, New York (1967), p. 42.
19 R. Bieber and G. Trumpler, Helv. Chim. Acta, 30, 1860
(1947).
20 L. C. Gruen and P. T. McTigue, J. Chem. Soc., 1963, 5217.
21 S. Kandlikar, B. Sethuram, and T. N. Rao, Indian J. Chem.,
Sect. A, 17, 264 (1979).
22 A. L. Jain and K. K. Banerji, J. Chem. Res. (M), 1983, 678.
23 K. K. Banerji, Tetrahedron, 43, 5949 (1987).
24 V. K. Sharma, K. Sharma, and N. Mishra, Oxid. Commun.,
16, 33 (1993).
25 S. Clementi, P. Linda, and G. Marino, Tetrahedron Lett.,
No. 17, 1389 (1970).
26 K. B. Wiberg and T. Mill, J. Am. Chem. Soc., 80, 3022
(1958).
27 K. B. Wiberg and W. H. Richardson, J. Am. Chem. Soc., 84,
2800 (1962).
28 J. Rocek, Tetrahedron Lett., No. 5, 1 (1959).
29 E. Lucchi, Gazz. Chim. Ital., 71, 729 (1941).
30 G. T. E. Graham and F.H. Westheimer, J. Am. Chem. Soc.,
80, 3030 (1958).
31 J. Rocek, F. H. Westheimer, A. Eschenmoser, L.
Moldovanyi, and J. Schreiber, Helv. Chim. Acta, 45, 2554 (1962).
32 J. Rocek and C. S. Ng, J. Am. Chem. Soc., 96, 1522 (1974).
33 U. Klanning, Acta Chem. Scand., 11, 313 (1957) ; 12, 576
(1958).
1
F. Freeman, J. B. Brant, N. B. Hester, A. A. Kamego, M. L.
Kasner, T. G. McLaughlin, and E. W. Paul, J. Org. Chem., 35, 982
(1970).
2
R. Gopalan and E. Kannamma, Indian J. Chem., Sect. A,
23, 518 (1984).
3
D. Kumar, A. Rani, D. N. S Prasad, and K. S Gupta, React.
Kinet. Catal. Lett., 43, 133 (1991).
4
(1998).
5
T. Veeraiah and S. Sondu, Indian J. Chem., Sect. A, 37, 328
K. G. Sekar, B. Ramkumar, and R. Rajaji, Oxid. Commun.,
24, 364 (2001).
6
(2001).
7
K. G. Sekar and M. Ravishankar, Oxid. Commun., 24, 368
B. Kuotsu, E. Tiewsoh, A. Debroy, and M. K. Mahanti, J.
Org. Chem., 61, 8875 (1996), and references therein; R.
Kharmutee, A. Debroy, and M. K. Mahanti, Oxid. Commun., 21,
553 (1998); E. Karim and M. K. Mahanti, Oxid. Commun., 21,
559 (1998); N. Thangkhiew, A. Debroy, and M. K. Mahanti, Oxid.
Commun., 22, 136 (1999).
8
M. K. Mahanti, Oxid. Commun., 22, 142 (1999); G. S.
Chaubey and M. K. Mahanti, Oxid. Commun., 23, 500 (2000); G.
G. Kharnaior, G. S. Chaubey, and M. K. Mahanti, Oxid.
Commun., 24, 377 (2001).
9
K. Balasubramanian and V. Prathiba, Indian J. Chem., Sect.
B, 25, 326 (1986).
10 R. P. Bell, Adv. Phys. Org. Chem., 4, 1 (1964).
11 K. B. Wiberg, “Oxidation in Organic Chemistry (Part A),”
Academic Press, New York (1965), p. 69.
12 K. K. Banerji, Indian J. Chem., Sect. A, 17, 300 (1979).
13 “Handbook of Chemistry and Physics,” ed by R. C. Weast,
CRC Press, Ohio (1978).
14 J. S. Littler and W. A. Waters, J. Chem. Soc., 1959, 1299.
15 N. Bailey, A. Carrington, K. A. K. Lott, and M. C. R.
Symons, J. Chem. Soc., 1960, 290; Y. Sasaki, Acta Chem. Scand.,
16, 719 (1962).
16 G. Michel and R. Machiroux, J. Raman Spectrosc., 14, 22
(1983); G. Michel and R. Cahay, J. Raman Spectrosc., 17, 79
34 F. A. Cotton and G. Wilkinson, “Advanced Inorganic
Chemistry,” Wiley Eastern, New Delhi (1985), p. 115.
35 C. G. Swain, R. F. W. Bader, R. M. Estene, and R. N.
Griffin, J. Am. Chem. Soc., 83, 1951 (1961).
36 T. R. Arnold and M. J. Danzig, J. Am. Chem. Soc., 79, 892
(1957).
37 J. S. Littler, Tetrahedron, 27, 81 (1971).
38 F. H. Westheimer, Chem. Rev., 45, 419 (1949).
39 J. F. Perez-Benito, C. Arias, and D. Lamrhari, Chem.
Commun., 1992, 472.