Kinetics and mechanism of the oxidation of aliphatic alcohols by quinolinium fluorochromate

Article abstract of DOI:10.1002/(SICI)1097-4601(1999)31:7<469::AID-KIN1>3.0.CO;2-2

The oxidation of some aliphatic alcohols by quinolinium fluorochromate (QFC) in dimethyl sulfoxide leads to the formation of corresponding carbonyl compounds. The reaction is first order with respect to QFC. The reaction exhibited Michaelis-Menten type kinetics with respect to the alcohol. The reaction is catalyzed by hydrogen ions. The hydrogen-ion dependence has the form: k_obs = a+b[H⁺]. The oxidation of [1,1-²H₂]ethanol (MeCD₂OH) exhibits a substantial primary kinetic isotope effect. The reaction has been studied in nineteen different organic solvents. The solvent effect was analyzed using Taft's and Swain's multiparametric equations. The rate of disproportionation of the complex is susceptible to both polar and steric effects of the substituents. A suitable mechanism has been proposed.

Full text of DOI:10.1002/(SICI)1097-4601(1999)31:7<469::AID-KIN1>3.0.CO;2-2

Kinetics and Mechanism of
the Oxidation of Aliphatic
Alcohols by Quinolinium
Fluorochromate
K. CHOUDHARY, P. K. SHARMA, K. K. BANERJI
Department of Chemistry, J.N.V. University, Jodhpur 342 005, India
Received 15 April 1998; accepted 26 January 1999
ABSTRACT: The oxidation of some aliphatic alcohols by quinolinium fluorochromate (QFC) in
dimethyl sulfoxide leads to the formation of corresponding carbonyl compounds. The reaction
is first order with respect to QFC. The reaction exhibited Michaelis-Menten type kinetics with
respect to the alcohol. The reaction is catalyzed by hydrogen ions. The hydrogen-ion depen-
dence has the form: kobs ϭ a ϩ b[Hϩ]. The oxidation of [1,1-2H2]ethanol (MeCD2OH) exhibits
a substantial primary kinetic isotope effect. The reaction has been studied in nineteen different
organic solvents. The solvent effect was analyzed using Taft’s and Swain’s multiparametric
equations. The rate of disproportionation of the complex is susceptible to both polar and
steric effects of the substituents. A suitable mechanism has been proposed. ᭧ 1999 John Wiley
& Sons, Inc. Int J Chem Kinet 31: 469–475, 1999
INTRODUCTION
EXPERIMENTAL
Materials
Quinolinium fluorochromate (QFC) has been used as
a mild and selective oxidizing reagent in synthetic or-
ganic chemistry [1]. However, there does not seem to
be many reports on the kinetic and mechanistic aspects
of oxidation reactions of QFC [2]. The kinetics of ox-
idation of alcohols by QFC has not been reported. We
have been interested in the kinetics of reactions of
complexed Cr(VI) species and have reported kinetics
and mechanism of the oxidation of aliphatic alcohols
by pyridinium fluorochromate (PFC) [3], pyridinium
chlorochromate (PCC) [4] and 2,2’-bipyridinium chlo-
rochromate(BPCC) [5]. It was observed that the oxi-
dations by PCC and PFC presented different kinetic
pictures. In this paper we describe the kinetics and
mechanism of the oxidation of nine primary and five
secondary aliphatic alcohols by QFC in dimethyl sulf-
oxide (DMSO) as solvent.
The QFC was prepared by the reported method [1]
and its purity was checked by an iodometric method.
The procedure used for the purification of alcohols
have been described earlier [6]. [1,1-2H2]Ethanol
(MeCD2OH) was prepared by Kalpan’s method [7].
Its isotopic purity, as ascertained by its NMR spectra,
was 93 Ϯ 5%. Due to the non-aqueous nature of the
medium, p-toluenesulphonic acid (TsOH) was used
as a source of hydrogen ions. TsOH is a strong acid
and in a polar medium like DMSO it is likely to be
completely ionised. Solvents were purified by the
usual method [8].
Product Analysis
The product analysis were carried out under kinetic
conditions. In a typical experiment, ethanol (2.30 g,
Correspondence to: K. K. Banerji
᭧ 1999 John Wiley & Sons, Inc.
CCC 0538-8066/99/070469-07

470
CHOUDHARY, SHARMA, AND BANERJI
0.05 mol) and QFC (2.49 g, 0.01 mol) were made up
to 50 cm3 in DMSO and kept in dark for about 15
hours to ensure the completion of the reaction. The
solution was then treated with an excess (200 cm3) of
a saturated solution of 2,4-dinitophenylhydrazine in
R2CHOH ϩ CrO2FOϪQHϩ !:
R2CO ϩ H2O ϩ CrOFOϪOHϩ (1)
Thus QFC undergoes a two-electron change. This ac-
cords with the earlier observations with PFC [9] and
QFC [2].
3
2 mol dmϪ HCl and kept overnight in a refrigerator.
The precipitated 2,4-dinitrophenylhydrazone (DNP)
was filtered off, dried, weighed, recrystalized from
ethanol and weighed again. The yield of DNP before
and after recrystallization was 2.0 g (90%) and 1.7 g
(75%), respectively. The DNP was found identical
(m.p. and mixed m.p.) with the DNP of acetaldehyde.
Similar experiments with other alcohols led to the for-
mation of DNP of the corresponding carbonyl com-
pounds in yields ranging from 71 to 83%, after re-
crystallization. Iodometric determinations of the
oxidation state of chromium in completely reduced re-
action mixtures indicated that the oxidation state of
the reduced chromium species was 3.92 Ϯ 0.21.
Rate Law
The reaction is first order with respect to QFC. Mi-
chaelis-Menten type kinetics were observed with re-
spect to alcohols (Table I). A plot of 1/kobs vs 1/[al-
cohol] is linear with an intercept on the rate ordinate.
This indicates the following overall mechanism equa-
tions (2) and (3) and the rate law equation (4).
Alcohol ϩ QFC IK
R
J [Complex]
(2)
(3)
k2
[Complex] 99: Product
Spectral Studies
Rate ϭ k2 K [Alcohol] [QFC]/
UV-VIS spectra (200–800 nm) of (i) QFC (0.001 mol
(1 ϩ K [Alcohol]) (4)
1
1
dmϪ ) alone and (ii) QFC (0.001 mol dmϪ ) ϩ
1
2-propanol (0.01 mol dmϪ ), in DMSO as the solvent,
The variation in alcohol concentration was studied at
four temperatures and the values of k2 and K were
evaluated from the double reciprocal plots. The ther-
modynamic parameters for the complex formation and
the activation parameters for its disproportionation
were also calculated (Tables II and III).
were obtained on a HP-diode array fast scanning spec-
trophometer (Model 8452A) at room temperature
(Ϸ300 K). The blank was DMSO. The scanning speed
1
was 600 nm sϪ . The time elapsed between the mixing
of the two reagents and recording of the spectrum (ii)
was Ͼ20 s.
Kinetic Measurements
Table I Rate Constants for the Oxidation of Ethanol
by QFC at 298 K
The reactions were followed under pseudo-first-order
conditions by keeping a large excess (ϫ15 or greater)
of the alcohol over QFC. The temperature was kept
constant to Ϯ0.1 K. The solvent was DMSO, unless
specified otherwise. The reactions were followed by
monitoring the decrease in the concentration of QFC
spectrophotometrically at 356 nm for 80% of the re-
action. The pseudo-first-order rate constants, kobs, were
evaluated from the linear (r ϭ 0.990–0.999) plots
of log [QFC] against time. Duplicate kinetic runs
showed that the rate constants were reproducible to
within Ϯ3%. Multivariate and simple linear regres-
sion analyses were carried out by the least-squares
method.
103 [QFC]
[Ethanol]
[TsOH]
104 kobs
Ϫ3
Ϫ3
Ϫ
3
Ϫ1
(mol dm
)
(mol dm
)
(mol dm
)
(s )
1.0
1.0
1.0
1.0
1.0
1.0
1.0
2.0
4.0
6.0
8.0
1.0
1.0
1.0
1.0
1.0
1.0
0.02
0.04
0.08
0.20
0.40
0.80
1.00
0.02
0.02
0.02
0.02
0.02
0.04
0.04
0.04
0.04
0.04
0.0
0.0
0.0
0.0
0.0
0.0
0.0
0.0
0.0
0.0
0.0
0.0
0.1
0.3
0.5
0.8
1.0
10.9
20.4
35.9
66.6
93.0
116
122
11.4
10.3
10.0
11.8
11.5*
23.1
28.2
34.4
43.2
50.1
RESULTS
The oxidation of alcohols results in the formation of
corresponding carbonyl compounds. The overall re-
action may be represented as equation (1).
* contained 0.001 mol dmϪ3 acrylonitrile

OXIDATION OF ALIPHATIC ALCOHOLS
471
Table II Formation Constants and Thermodynamic Parameters of Alcohol-QFC Complexes in DMSO
Ϫ1
K/dm3 mol
⌬ H
⌬ S
Ϫ
1
Ϫ
Ϫ1
S. No.
Alcohol
Methanol
Ethanol
Propanol-1
Butanol-1
Pentanol-1
2-Methyl-propanol-1
2-Chloro-ethanol
2-Methoxy-ethanol
2,2-Dimethyl-propanol-1
Propanol-2
Butanol-2
Pentanol-2
1-Chloro-propanol-2
1-Methoxy-propanol-2
MeCD2OH
288 K
298 K
308 K
318 K
(kJ mol
)
(J mol 1K
)
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
3.52
4.51
5.73
4.10
4.75
5.15
4.80
3.95
4.85
5.32
3.91
4.39
5.10
4.50
4.50
2.85
3.81
4.91
3.31
3.90
4.45
4.15
3.15
4.21
4.50
3.35
3.63
4.50
3.63
3.61
2.33
3.30
4.11
2.69
3.21
3.81
3.41
2.50
3.53
3.80
2.50
2.86
3.73
2.89
2.90
1.68
2.65
3.20
2.00
2.45
3.10
2.71
1.91
2.85
2.83
1.82
2.08
3.03
2.40
2.11
Ϫ21.2 Ϯ 0.8
Ϫ15.7 Ϯ 0.7
Ϫ17.1 Ϯ 0.9
Ϫ20.5 Ϯ 0.9
Ϫ19.1 Ϯ 0.8
Ϫ15.2 Ϯ 0.7
Ϫ17.0 Ϯ 0.9
Ϫ20.7 Ϯ 0.8
Ϫ15.9 Ϯ 0.8
Ϫ18.1 Ϯ 1.0
Ϫ22.1 Ϯ 1.1
Ϫ21.3 Ϯ 1.0
Ϫ15.8 Ϯ 0.9
Ϫ18.6 Ϯ 0.8
Ϫ21.9 Ϯ 0.6
Ϫ55 Ϯ 3
Ϫ34 Ϯ 2
Ϫ36 Ϯ 3
Ϫ51 Ϯ 3
Ϫ45 Ϯ 3
Ϫ31 Ϯ 2
Ϫ38 Ϯ 3
Ϫ52 Ϯ 3
Ϫ34 Ϯ 2
Ϫ41 Ϯ 3
Ϫ57 Ϯ 4
Ϫ53 Ϯ 3
Ϫ33 Ϯ 2
Ϫ44 Ϯ 2
Ϫ55 Ϯ 2
Induced Polymerization of Acrylonitrile
2H2]ethanol was studied. The results showed the pres-
ence of a substantial primary kinetic isotope effect
(Table III).
The oxidation of alcohols, in an atmosphere of nitro-
gen, failed to induce polymerization of acrylonitrile.
Further, the addition of acrylonitrile did not affect the
rate. This indicates that a one-electron oxidation, giv-
ing rise to free radicals, is unlikely in the present re-
action (Table I).
Spectral Studies
A perusal of the spectra of QFC and QFC ϩ
2-propanol showed that there is a distinct change in
the absorption spectra of QFC on the addition of 2-
propanol (Figure 1). It was further observed that the
absorbance due to the intermediate decreases with
time. This supports the postulation of the formation of
Kinetic Isotope Effect
To ascertain the importance of cleavage of the C9H
bond in the rate-determining step, oxidation of [1,1-
Table III Rate Constants of Disproportionation and Activation Parameters of the Alcohol-QFC Complexes in DMSO
104 k2/s
298 K
2.65
154
256
452
510
765
3.8
28.2
7660
8890
14000
25000
382
⌬ H*
⌬ S*
⌬ G*
Ϫ
1
Ϫ
1
Ϫ
Ϫ1
Ϫ1
Alcohola
288 K
0.95
71.0
116
308 K
7.68
318 K
(kJ mol
)
(J mol 1K
)
(kJ mol )
1
2
3
21.8
748
78.1 Ϯ 0.9
57.4 Ϯ 0.8
56.7 Ϯ 0.5
52.8 Ϯ 0.8
51.8 Ϯ 0.9
49.9 Ϯ 0.9
73.7 Ϯ 0.9
64.1 Ϯ 0.8
47.8 Ϯ 0.2
47.8 Ϯ 1.1
46.6 Ϯ 1.0
49.8 Ϯ 1.2
64.2 Ϯ 0.7
56.7 Ϯ 1.0
62.1 Ϯ 0.8
Ϫ52 Ϯ 3
Ϫ87 Ϯ 3
Ϫ86 Ϯ 2
Ϫ94 Ϯ 3
Ϫ96 Ϯ 3
Ϫ99 Ϯ 3
Ϫ63 Ϯ 3
Ϫ79 Ϯ 2
Ϫ87 Ϯ 1
Ϫ86 Ϯ 3
Ϫ86 Ϯ 3
Ϫ71 Ϯ 3
Ϫ57 Ϯ 2
Ϫ68 Ϯ 3
Ϫ86 Ϯ 3
93.3 Ϯ 0.7
83.3 Ϯ 0.7
82.1 Ϯ 0.4
80.6 Ϯ 0.7
80.4 Ϯ 0.7
79.4 Ϯ 0.8
92.4 Ϯ 0.8
87.5 Ϯ 0.6
73.7 Ϯ 0.1
73.3 Ϯ 1.0
72.1 Ϯ 1.0
70.8 Ϯ 1.4
81.1 Ϯ 0.9
76.9 Ϯ 1.1
87.5 Ϯ 0.9
345
565
940
1190
1950
2130
2980
27.5
4
5
6
220
248
370
1020
1440
7
8
1.40
11.8
11.2
71.5
14600
17100
26700
48900
932
160
9
3750
4400
7100
11800
156
925
12.9
5.50
27400
32000
49200
93000
2150
9500
164
10
11
12
13
14
15
kH/kD
2100
30.1
5.12
4610
71.5
4.83
4.56
a For the identity of compounds see Table II

472
CHOUDHARY, SHARMA, AND BANERJI
Solvent Effect
The oxidation of ethanol was studied in 19 different
organic solvents. The choice of solvent was limited
due to the solubility of QFC and its reaction with pri-
mary and secondary alcohols. There was no reaction
with the solvents chosen. The kinetics were similar in
all the solvents. The values of K and k2 are recorded
in Table IV.
DISCUSSION
There is no significant isokinetic relationship between
the activation enthalpies and entropies of the oxidation
of fourteen alcohols (r2 ϭ 0.5112). A correlation be-
tween the calculated values of enthalpies and entropies
of activation is often vitiated by the experimental er-
rors associated with them. Exner [10] has suggested
an alternative method of testing the validity of isoki-
netic relationship. An Exner’s plot between log k2 at
288 K and at 318 K was linear (r2 ϭ 0.9970). The
value of isokinetic temperature evaluated from the
Exner’s plot is 1510 Ϯ 62K. The linear isokinetic
correlation implies that all the alcohols are oxidized
by the same mechanism and are governed by the
changes in both the enthalpy and entropy of the acti-
vation.
Ϫ
Figure 1 UV-VIS spectra of (A) 0.001 mol dm 3 QFC and
(B) 0.001 mol dm 3 QFC ϩ 0.10 mol dm 3 2-propanol. Sol-
Ϫ
Ϫ
vent : DMSO
an intermediate in a rapid pre-equilibrium and its sub-
sequent decomposition to yield the ultimate product.
Effect of Acidity
The reaction is catalyzed by hydrogen ions. The hy-
drogen-ion dependence has the following form equa-
tion (5). The results are recorded in Table I.
kobs ϭ a ϩ b [Hϩ]
(5)
Solvent Effect
To determine whether the acidity is affecting the
formation constant, K, and/or the rate constant, k2, of
the disproportionation, the effect of variation in the
alcohol concentration was studied at different acidi-
ties. Results showed that K is not appreciably affected
by changes in the acidity whereas the value of k2 vary.
The values of a and b, for ethanol, are 1.96 Ϯ 0.1 ϫ
The rate constants, k2, for the disproportionation of the
complex, in eighteen solvents (CS2 was not consid-
ered, as the complete range of solvent parameters was
not available) were correlated in terms of the linear
solvation energy relationship equation (7) of Kamlet
et al. [12].
3
1
3
1
1
10Ϫ sϪ and 3.00 Ϯ 0.1 ϫ 10Ϫ molϪ dm3 sϪ re-
spectively (r2 ϭ 0.9986).
log k2 ϭ A0 ϩ p␲* ϩ b␤ ϩ a␣
(7)
Table IV Formation Constants and Rate Constants for the Decomposition of Ethanol-QFC Complexes in Different
Solvents at 298 K
K
103 k2
K
103 k2
Ϫ1
Ϫ
1
Ϫ1
Ϫ1
Solvents
Chloroform
1,2-Dichloromethane
Dichloromethane
DMSO
Acetone
N,N-Dimethylformamide
Butanone
dm mol
s
Solvents
Acetic acid
Cyclohexane
Toluene
Acetophenone
THF
t-butyl alcohol
1,4-Dioxane
1,2-Dimethoxyethane
Carbon disulfide
dm mol
s
6.21
5.76
4.90
4.51
7.35
7.14
6.45
5.86
8.11
4.90
18.6
22.9
20.4
71.0
19.1
37.2
7.24
24.5
6.92
7.59
7.10
7.12
7.63
4.80
6.57
8.00
7.21
6.55
5.85
3.02
0.59
5.01
28.2
9.33
7.24
10.2
4.79
2.45
Nitrobenzene
Benzene
Ethyl acetate

OXIDATION OF ALIPHATIC ALCOHOLS
473
In this equation, ␲* represents the solvent polarity, ␤
the hydrogen bond acceptor basicities, and ␣ is the
hydrogen bond donor acidity. A0 is the intercept term.
It may be mentioned here that out of the 18 solvents,
13 have a value of zero for ␣. The results of correlation
analyses in terms of eq. (7), a biparametric equation
involving ␲* and ␤, and separately with ␲* and ␤ are
given below in equations. (8)–(11).
log k2
ϭ 0.65(Ϯ0.02) A ϩ 1.85 (Ϯ0.01) B Ϫ 4.36 (13)
R2 ϭ 0.9990; sd ϭ 0.02; n ϭ 19; ␥ ϭ 0.02
log k2 ϭ 0.39(Ϯ0.61) A Ϫ 3.10
r2 ϭ 0.0236; sd ϭ 0.49; n ϭ 19; ␥ ϭ 0.95
log k2 ϭ 1.79(Ϯ0.12) B Ϫ 4.15
(14)
(15)
log k2
ϭ Ϫ4.12 ϩ 1.71 ␲* ϩ 0.21 ␤ ϩ 0.16␣
(Ϯ0.22) (Ϯ0.18) (Ϯ0.17) (8)
r2 ϭ 0.9336; sd ϭ 0.13; n ϭ 19; ␥ ϭ 0.19
R2 ϭ 0.8527; sd ϭ 0.20; n ϭ 18; ␥ ϭ 0.30
log k2 ϭ 1.45 Ϯ 0.15 (A ϩ B) Ϫ 4.33 (16)
r2 ϭ 0.8582; sd ϭ 0.17; n ϭ 19; ␥ ϭ 0.28
log k2 ϭ Ϫ4.16 ϩ 1.77 (Ϯ0.21) ␲*
ϩ 0.15 (Ϯ0.17) ␤ (9)
The rates of disproportionation of the complex in
different solvents showed an excellent correlation in
Swain’s equation with both cation and anion-solvating
powers contributing toward the observed effect of the
solvents. However, the contribution of the cation-sol-
vation is relatively greater. A correlation with B alone
showed that the cation-solvation accounts for about
93% of the data. The solvent polarity, represented by
(A ϩ B), accounted for about 86% of the data. In view
of the fact that solvent polarity is able to account for
about 86% of the data, an attempt was made to cor-
relate the rate with the relative permittivity of the sol-
vent. It was observed that a plot of log (rate) against
the inverse of the relative permittivity is not linear
(r2 ϭ 0.5150; sd ϭ 0.35; ␥ ϭ 0.55).
R2 ϭ 0.8431; sd ϭ 0.19; n ϭ 18; ␥ ϭ 0.30
log k2 ϭ Ϫ4.13 ϩ 1.81 (Ϯ0.20) ␲*
r2 ϭ 0.8348; sd ϭ 0.20; n ϭ 18; ␥ ϭ 0.30
log k2 ϭ Ϫ3.13 ϩ 0.47 (Ϯ0.39) ␤
(10)
(11)
r2 ϭ 0.829; sd ϭ 0.47; n ϭ 18; ␥ ϭ 0.87
Here n is the number of data points and ␥ is the Ex-
ner’s statistical parameter [13].
Kamlet’s [12] triparametric equation explains about
85% of the effect of solvent on the oxidation. How-
ever, by Exner’s criterion [13] the correlation is not
even satisfactory (cf. eq. 8). The major contribution
is of solvent polarity. It alone accounted for about
83% of the data. Both ␤ and ␣ play relatively minor
roles.
Correlation Analysis of Reactivity
The rates of oxidation of the alcohols failed to yield
any significant correlation separately with Taft’s ␴*
and Es values equations (17) and (18)].
The data on the solvent effect were analyzed in
terms of Swain’s equation [14] of cation- and anion-
solvating concept of the solvents as well as equation
(12).
log k2 ϭ Ϫ2.13(Ϯ0.34) ͚ ␴* 0.09
r2 ϭ 0.7669; sd ϭ 0.53; ␥ ϭ 0.37; n ϭ 14
log k2 ϭ Ϫ1.22(Ϯ0.38) ͚ Es Ϫ 0.25
r2 ϭ 0.7111; sd ϭ 0.66; ␥ ϭ 0.41; n ϭ 14
(17)
(18)
log k2 ϭ aA ϩ bB ϩ C
(12)
Here A represents the anion-solvating power of the
solvent and B the cation-solvating power. C is the in-
tercept term. (A ϩ B) is postulated to represent the
solvent polarity. The rates in different solvents were
analyzed in terms of equation (12), separately with A
and B and with (A ϩ B).
The rates were, therefore, correlated in terms of Pav-
elich-Taft’s [15] dual substituent-parameter equation
(19).
log k2 ϭ * ͚ ␴* ϩ ␦ ͚ Es ϩ log k0
(19)

474
CHOUDHARY, SHARMA, AND BANERJI
Table V Temperature Dependence of the Reaction Constants
Temp./K
␳*
␦
R2
sd
␥
288
298
308
318
Ϫ1.68 Ϯ 0.04
Ϫ1.60 Ϯ 0.03
Ϫ1.49 Ϯ 0.04
Ϫ1.42 Ϯ 0.05
Ϫ0.82 Ϯ 0.02
Ϫ0.81 Ϯ 0.02
Ϫ0.79 Ϯ 0.02
Ϫ0.76 Ϯ 0.03
0.9983
0.9984
0.9981
0.9966
0.06
0.05
0.06
0.07
0.03
0.03
0.03
0.04
The values of substituent constants were obtained
from the compilation by Wiberg [16]. The correlations
are excellent: the reaction constants being negative
(Table V). There is no significant collinearity (r2 ϭ
0.2322) between ␴* and Es values of the fourteen sub-
stituents.
The negative polar reaction constant indicates an
electron-deficient carbon center in the transition state
of the rate-determining step. The negative steric re-
action constant shows a steric acceleration of the re-
action. This may be explained on the basis of high
ground state energy of the sterically crowded alcohols.
Since the crowding is relieved in the product carbonyl
compound as well as in the transition state leading to
it, the transition state energies of the crowded and un-
crowded alcohols do not differ much and steric accel-
eration, therefore, results.
polar reaction constant together with substantial deu-
terium isotope effect indicate that the transition state
approaches a carbocation in character. Hence, the
transfer of hydride-ion from alcohol to the oxidant is
suggested. The hydride-transfer mechanism is also
supported by the major role of cation-solvating power
of the solvents. The kinetics and spectral studies point
to the formation of an intermediate in the pre-equilib-
rium. Therefore, it is likely that the oxidation involves
a hydride-ion transfer via a chromate ester. This pos-
tulation is supported by an analysis of the temperature
dependence of kinetic isotope effect.
Kwart and Nickle [17] have shown that a study of
the dependence of the kinetic isotope effect on tem-
perature can be gainfully employed to resolve this
problem. The data for protio- and deuterio-ethanols,
fitted to the familiar expression kH/kD ϭ AH/AD
exp(Ea/RT) [18,19] show a direct correspondence with
the properties of a symmetrical transition state in
which the activation energy difference (Ea) for kH/kD
is equal to the zero-point energy difference for the
respective C9H and C9D bonds (Ϸ4.5 kJ/mol) and
the frequency factors and the entropies of activation
Mechanism
The presence of a substantial primary kinetic isotope
effect confirms the cleavage of ␣9C9H bond in the
rate-determining step. The large negative value of the
ACID-DEPENDENT PATH
ACID-INDEPENDENT PATH
O
OϪQHϩ
OH
K
ϩ
R29C9OH ϩ
Cr
(A) ϩ [Hϩ]
R29C9O9Cr9OϪQHϩ
O
F
H
H
OH
F
O
k2
R29C9O9Cr9OϪQHϩ
#
OH
␦ϩ
Cr9OϪQHϩ
H
OH
F
␦ϩ
R29C
O
(A)
k2
H
OH
F
#
O
ϩ
HOCrFOQϪHϩ ϩ H2O ϩ R29CO
R29C
O
Cr9OϪQHϩ
H
OH
F
OCrFOQϪHϩ ϩ H2O ϩ R29CO
Scheme I

OXIDATION OF ALIPHATIC ALCOHOLS
475
3. Banerji, K. K. J Chem Soc, Perkin Trans 1988, 2, 547.
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5. Rathore, S.; Sharma, P. K.; Banerji, K. K. Indian J Chem
1995, 34B, 702.
of the respective reactions are nearly equal. The sim-
ilar phenomena were observed earlier in the oxidation
of diols by BPCC [20] and that of alcohols by PFC
[3].
6. Mathur, D.; Sharma, P. K.; Banerji, K. K. J Chem Soc,
Perkin Trans 1993, 2, 205. Grover, A.; Varshney, S.;
Banerji, K. K. Indian J Chem, Sect A 1996, 35, 206.
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tion of organic compounds; Pergamon Press: Oxford,
1966.
9. Banerji, K. K. J Chem Soc, Perkin Trans 1988, 2, 2065.
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1094.
11. Khanchandani, R.; Sharma, P. K.; Banerji, K. K. J Chem
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R. W. J Org Chem 1983, 48, 2877 and references cited
therein.
13. Exner, O. Collect Czech Chem Commun 1966, 31,
3222.
14. Swain, C. G.; Swain, M. S.; Powel, A. L.; Alunni, S. J
Am Chem Soc 1983, 105, 502.
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York, 1963; p. 416.
Bordwell [21] has documented a very cogent evi-
dence against the occurrence of concerted one-step
biomolecular processes by hydrogen transfer and it is
evident that in the present studies also the hydrogen
transfer does not occur by an acyclic biomolecular
process. It is well established that intrinsically con-
certed sigmatropic reactions, characterized by transfer
of hydrogen in a cyclic transition state, are the only
truly symmetrical processes involving a linear hydro-
gen transfer [22]. Littler [23] has also shown that a
cyclic hydride transfer, in the oxidation of alcohols by
Cr(VI), involves six electrons and, being a Huckel-
type system, is an allowed process. Thus, the overall
mechanism is proposed to involve the formation of a
chromate ester in a fast pre-equilibrium step and then
a disproportionation of the ester in a subsequent slow
step via a cyclic concerted symmetrical transition state
leading to the product (Scheme I). The observed hy-
drogen-ion dependence can be explained by assuming
a rapid reversible protonation of the chromate ester (A)
with the protonated ester decomposing at a rate faster
than (A).
17. Kwart, H.; Nickel, J. H. J Am Chem Soc 1973, 95,
3394.
Thanks are due to the University Grants Commission (India)
for the financial support.
18. Kwart, H.; Latimer, M. C. J Am Chem Soc 1971, 93,
3770.
19. Kwart, H.; Slutsky, J. J Chem Soc, Chem Commun
1972, 1182.
BIBLIOGRAPHY
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1997, (S), 242; (M), 1663.
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Engl 1969, 8, 781.
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