Kinetic and mechanistic aspects for the tartaric acid oxidation by vanadium(V) in sulfuric acid medium

Article abstract of DOI:10.1002/(SICI)1097-4601(1998)30:1<55::AID-KIN7>3.0.CO;2-T

Tartaric acid oxidation by vanadium(V) in sulfuric acid medium was investigated spectrophotometrically at 760 nm and 30°C by appearance of the vanadium(IV), as vanadyl. The reaction rate was determined under pseudo-first-order conditions with an excess of hydroxyacid over the oxidant concentration. The oxidation showed a first-order dependence with respect to vanadium(V) concentration and fractional orders with respect to tartaric acid and sulfuric acid concentrations, with no control and with constant ionic strength, The reaction rate is enhanced by an increase of ionic strength, and slightly reduced by a decrease of the dielectric constant of the medium. The activation parameters were calculated based on the rate constants determined in the 293 to 313 K interval. The proposed oxidation mechanisms and the derived rate laws are consistent with the experimental rate laws.

Full text of DOI:10.1002/(SICI)1097-4601(1998)30:1<55::AID-KIN7>3.0.CO;2-T

Kinetic and Mechanistic
Aspects for the Tartaric
Acid Oxidation by
Vanadium(V) in Sulfuric
Acid Medium
´
´
KEIKO TAKASHIMA, CLAUDIO M. ZIGLIO, CELIA M. RONCONI
Departamento de Quı´mica, Universidade Estadual de Londrina, caixa postal 6001, 86051-970, Londrina, PR, Brasil
Received 7 October 1996; accepted 13 May 1997
ABSTRACT: Tartaric acid oxidation by vanadium(V) in sulfuric acid medium was investigated
spectrophotometrically at 760 nm and 30ЊC by appearance of the vanadium(IV), as vanadyl.
The reaction rate was determined under pseudo-first-order conditions with an excess of hy-
droxyacid over the oxidant concentration. The oxidation showed a first-order dependence with
respect to vanadium(V) concentration and fractional orders with respect to tartaric acid and
sulfuric acid concentrations, with no control and with constant ionic strength. The reaction
rate is enhanced by an increase of ionic strength, and slightly reduced by a decrease of the
dielectric constant of the medium. The activation parameters were calculated based on the
rate constants determined in the 293 to 313 K interval. The proposed oxidation mechanisms
and the derived rate laws are consistent with the experimental rate laws. ᭧ 1998 John Wiley &
Sons, Inc. Int J Chem Kinet: 30: 55–61, 1998.
INTRODUCTION
an ionic mechanism involving transfer of an ␣-hydro-
gen or by concerted oxidative decarboxylation involv-
ing cleavage of a C9C bond [7]. The choice of the
mechanism route would thus depend on several fac-
tors, such as the structure of the substrates and the
nature and the oxidizing capacity of the oxidant, the
reaction conditions, and the actual products isolated.
Nevertheless, relatively few studies have concerned
themselves with the nature of the oxidizing species in
these systems [8,9] because apparently the kinetic be-
havior of hydroxyacids oxidations are well-estab-
lished.
In our earlier study [10–12] we have reported that
different mechanisms are operative with the partici-
pation of some different active species of vana-
dium(V). The oxidation rate of malic acid by vana-
dium(V) was greater in sulfuric acid medium with
The kinetics of ␣-hydroxyacids oxidation through sev-
eral oxidants has been investigated [1–4] and recently
reviewed by Wells [5]. Being bifunctional com-
pounds, their oxidation can proceed by many possible
routes. While oxidation of hydroxyacids by one-elec-
tron oxidants generally take place through one of the
several possible initial radical formation steps [6], two
electrons oxidants usually promote oxidation either by
Correspondence to: K. Takashima
Contract grant Sponsor: PADCT/FINEP
Contract grant Sponsor: PADCT/CAPES
Contract grant Sponsor: CNPq
Contract grant Sponsor: CPG/UEL
International Journal of Chemical Kinetics, Vol. 30, 55–61 (1998)
᭧ 1998 John Wiley & Sons, Inc.
CCC 0538-8066/98/010055-07

56
TAKASHIMA, ZIGLIO, AND RONCONI
ionic strength control and showed first-order behavior
with respect to oxidant, substrate, and hydrogen ion
concentrations with or without control of the ionic
strength [10]. However, when this oxidation was per-
formed in perchloric acid medium, the rate constant
remained practically constant with increase of the
ionic strength. The rate law showed a first-order be-
havior with respect to the oxidant and malic acid con-
centrations and a fractional-order with respect to the
hydrogen ion concentration and a similar mechanism
was proposed with or without control of ionic strength
[12]. On the other hand, when lactic acid was oxidized
in a sulfuric acid medium, a larger rate was observed
with ionic strength control, and the plot of ln kobs vs.
ln [Hϩ] displayed an inflexion in both conditions, in-
dicating two different orders in the investigated hy-
drogen ion concentrations range [11]. The variation of
the reaction-order was more significant when the ionic
strength was not controlled, showing a behavior of ap-
proximately first- and second-orders, respectively. We
attributed these changes not only to the nature of the
vanadium(V), but also to the oxidizing capacity pro-
moted by the increase of the hydrogen ion concentra-
tion. In continuation with our studies, we report herein
the results of the oxidation of tartaric acid by vana-
dium(V) in sulfuric acid medium under variable and
constant ionic strength conditions.
phenylhydrazine, and the Tollens reagent displayed to
be positive for aldehyde [13]. The production of for-
maldehyde was verified by means of the Hantzsch re-
action through the addition of reagents containing ace-
tylacetone and ammonium salt [14]. The stoichiometry
was determined as described early [11]. A yield of 2.0
(Ϯ0.2) moles of CO2 for each mol of HTA was ob-
tained in a Warburg respirometer (B. Braun, model
V-85) at 313 K. The overall reaction for tartaric acid
oxidation may be represented by the following equa-
tion
HOOCCH(OH)CH(OH)COOH ϩ 2 V(V) !:
2 HCOH ϩ 2 CO2 ϩ 2 V(IV) ϩ 2 Hϩ
RESULTS
Effect of Reactants Concentration
An enhancement in the oxidation rate was observed
with an increase of the excess tartaric acid concentra-
tion, [HTA], when the initial concentrations of V(V),
[V(V)]0, and sulfuric acid, [H2SO4]0, were main-
tained constant at 303 K as shown in Table I (r Ն
0.9962 and s Յ 0.0797). This was independent of
whether the reaction took place at constant or with no
control of the ionic strength. Individual plots of ln
[V(V)] as a function of time for different initial con-
centrations of tartaric acid, [HTA]0, display satisfac-
tory linearity (r Ն 0.9988 and s Յ 2.20 ϫ 10Ϫ4) at
EXPERIMENTAL
The solution of vanadium(V), V(V), was prepared
and the ionic strength was controlled as described pre-
viously [10,11]. The tartaric acid, HTA, (Merck, pu-
rity 99.5%) was used after recrystallization with water.
The pseudo-first-order rate constants, kobs were cal-
culated from the slopes of the linear plots (correlation
coefficient, r Ն 0.9988, standard deviation, s Յ
2.2 ϫ 10Ϫ4) of ln[V(V)] against time, and were re-
producible within Ϯ5%. The kinetic measurements
were performed monitoring the appearance of vana-
dium(IV), V(IV), at 760 nm with a Varian spectro-
photometer model DMS-80 at 303.0 K (Ϯ0.1 K). The
measurements were carried out in duplicate and, when
necessary, in triplicate.
The oxidation of tartaric acid by V(V) induced the
polymerization of a solution of recrystallized acrylam-
ide (25%) under kinetic conditions. Control experi-
ments in which either V(V) or HTA were excluded
demonstrated the absence of polymerization.
The product analysis was carried out under kinetic
conditions. The presence of aldehyde and/or ketone
was characterized by the reaction with 2,4-dinitro-
Table I Pseudo-First-Order Rate Constants for the
Oxidation of Tartaric Acid by V(V) with No Control of
Ionic Strength (I), and at Constant Ionic Strength.a
2
[V(V) ϭ 1.00 ϫ 10Ϫ mol dmϪ3, [H2SO4] ϭ 0.15 mol
dmϪ3, and T ϭ 303 K
Ϫ
3
Ϫ
3
Ϫ
Ϫ1
I/mol dm
[HTA]/mol dm
kobs/10 3 s
b
b
b
b
b
0.50
0.75
1.00
1.25
1.50
0.50
0.75
1.00
1.25
1.50
1.03
1.30
1.51
1.74
1.96
2.25
2.87
3.34
3.89
4.24
1.10
1.10
1.10
1.10
1.10
a Constant ionic strength was achieved by using NaHSO4 as an
added electrolyte.
b Initial ionic strength is 0.16 mol dmϪ3, but changes in the
course of the reaction.

TARTARIC ACID OXIDATION
57
Table II Pseudo-First-Order Rate Constants for the
Reduction of Vanadium(V) at 303 K with No Control of
Ionic Strength (I), and at Constant Ionic Strength.a
Table III Effect of the Sulfuric Acid Concentration on
the Pseudo-First-Order Rate Constants at 303 K with
No Control of Ionic Strength (I), and at Constant Ionic
3
2
[HTA] ϭ 1.00 mol dmϪ3 and [H2SO4] ϭ 0.15 mol dmϪ
Strength.a [V(V)] ϭ 1.00 ϫ 10Ϫ mol dmϪ3 and [HTA] ϭ
3
1.00 mol dmϪ
Ϫ
3
Ϫ
Ϫ
3
Ϫ
Ϫ1
I/mol dm
[V(V)]/10 2 mol dm
kobs/10 3 s
Ϫ
3
Ϫ
3
Ϫ
Ϫ1
I/mol dm
[H2SO4]/mol dm
kobs/10 3 s
b
b
b
b
b
0.50
1.00
1.50
2.00
2.50
0.50
1.00
1.50
2.00
2.50
1.35
1.51
1.66
1.86
1.98
3.00
3.34
3.51
3.72
3.98
0.06
0.11
0.16
0.26
0.41
0.56
0.71
0.86
1.01
1.10
1.10
1.10
1.10
1.10
1.10
1.10
1.10
1.10
0.05
0.10
0.15
0.25
0.40
0.55
0.70
0.85
1.00
0.05
0.10
0.15
0.25
0.40
0.55
0.70
0.85
1.00
0.91
1.21
1.51
2.31
3.30
3.87
4.15
4.51
5.11
2.56
2.87
3.34
3.58
4.16
4.63
4.98
5.58
5.85
1.10
1.10
1.10
1.10
1.10
a Constant ionic strength was achieved by using NaHSO4 as an
added electrolyte.
b There is a small variation in the initial ionic strength to the
initial [V(V)], which changes in the course of the reaction.
303 K, and the order with respect to [HTA] is around
0.5 (r Ն 0.9902 and s Յ 0.0301). The variation of
2
3
[V(V)]0 from 0.50 ϫ 10Ϫ to 2.5 ϫ 10Ϫ2 mol dmϪ
a Constant ionic strength was achieved by using NaHSO4 as an
added electrolyte.
increased the rate constant by about 47% (r ϭ 0.9978
and s ϭ 0.0193) in solutions with no control of ionic
strength, and by about 33% (r ϭ 0.9993 and s ϭ
0.0166) at constant ionic strength as indicated in
Table II.
Effect of Solvent
The solvent composition was varied by adding meth-
anol to aqueous medium (0–40% v/v). The oxidation
rate decreased with the increase of methanol content
(Table V). Plots of ln kobs vs. reciprocal of dielectric
constant of the medium, 1/⑀ [17] resulted in straight
lines (r Ն 0.9065, s Յ 0.05045). Control experiments
showed a negligible oxidation of methanol by V(V).
Effect of Hydrogen Ion Concentration
Maintaining the same initial reactants concentrations,
[V(V)]0 and [HTA]0, the oxidation rate increased with
the increase of the hydrogen ion concentration, [Hϩ],
through the addition of sulfuric acid. The pseudo-first-
order rate constants at 303 K, are presented in Table
III. The reaction-order for the reaction without ionic
strength control was 0.60 (r ϭ 0.9946 and s ϭ 0.0699)
and with control, 0.28 (r ϭ 0.9959 and s ϭ 0.0439).
Table IV Effect of the Ionic Strength (I) on the
Pseudo-First-Order Rate Constants at 303 K. [V(V)] ϭ
1.00 ϫ 10Ϫ2 mol dmϪ3, [HTA] ϭ 1.00 mol dmϪ3, and
3
[H2SO4] ϭ 0.15 mol dmϪ
Ϫ
3
Ϫ3
Ϫ1
I/mol dm
kobs/10
s
Effect of Ionic Strength
0.16
0.20
0.50
0.80
1.10
1.40
1.70
2.00
1.51
1.54
2.23
2.75
3.34
3.88
4.86
5.86
The kinetic experiments were performed changing the
ionic strength by the addition of a concentrated solu-
tion of sodium bisulfate (5.0 mol dmϪ3) in the reactant
solution. The pseudo-first-order rate constant in-
creased by approximately four-fold in the range of
3
0.16 to 2.00 mol dmϪ (Table IV). The plot of log
kobs as a function of the square root of the ionic
strength at 303 K [15,16] gives a product of the ionic
charges, zϩzϪ of 0.56 (r ϭ 0.9976 and s ϭ 0.0164).
a Ionic strength was achieved by using NaHSO4 as an added
electrolyte.

58
TAKASHIMA, ZIGLIO, AND RONCONI
Table V Pseudo-First-Order Rate Constants for the
Tartaric Acid Oxidation in Binary Aqueous Mixtures of
Methanol with No Control and at Constant Ionic
3
Strength (I)a at 303 K. [V(V)] ϭ 1.00 ϫ 10Ϫ2 mol dmϪ
,
[HTA] ϭ 1.00 mol dmϪ3, and [H2SO4] ϭ 0.15 mol dmϪ
3
Ϫ
3
Ϫ
Ϫ1
I/mol dm
[MeOH]/%
kobs/10 3 s
b
b
b
b
b
b
0
20
25
30
35
40
0
20
25
30
35
40
1.51
1.36
1.26
1.26
1.30
1.09
3.34
3.06
2.79
2.84
2.84
2.70
1.10
1.10
1.10
1.10
1.10
1.10
Ϫ1
Figure 1 Plot of ln kЈ vs. T for [V(V)] ϭ 1.00 ϫ
Ϫ
2
Ϫ
Ϫ
10 mol dm 3, [H2SO4] ϭ 0.15 mol dm 3, and [HTA] ϭ
0.50–1.50 mol dm 3. (᭿) No control of the ionic strength
and (᭹) constant ionic strength (I ϭ 1.10 mol dm 3).
Ϫ
Ϫ
a Initial ionic strength is 0.16 mol dmϪ3, but changes in the
course of the reaction.
b Ionic strength was achieved by using NaHSO4 as an added
electrolyte.
perature. The activation parameters are displayed in
Table VII.
Effect of Temperature
The oxidation rate was measured by varying the tem-
perature from 293 to 313 K, maintaining constant all
the other experimental conditions. These results are
given in Table VI (r Ն 0.9988 and s Յ 2.20 ϫ
10Ϫ4) and the reaction order with respect to [HTA]
was remained around 0.5. The Arrhenius plots (Fig.
1, r Ն 0.9961 and s Յ 0.08132) were constructed from
ln kЈ (the apparent second-order rate constants, ob-
tained from Table VI) vs. the reciprocal absolute tem-
DISCUSSION
The reaction rate was found to be linearly proportional
to the concentration of the oxidant, and to have a frac-
tional dependence on the concentrations of both tar-
taric acid and hydrogen ion under constant and vari-
able ionic strengths. We assumed that sulfuric acid
Table VI Effect of Temperature on the Tartaric Acid Oxidation with No Control of Ionic Strength (I), and at
2
Constant Ionic Strength.a [V(V)] ϭ 1.00 ϫ 10Ϫ mol dmϪ3, [HTA] ϭ 0.50–1.50 mol dmϪ3, and [H2SO4] ϭ 0.15 mol
3
dmϪ
Ϫ
Ϫ1
kobs/10 3 s
Ϫ
3
Ϫ3
T/K
I/mol dm
[HTA]/mol dm
0.50
0.75
1.00
1.25
1.50
b
b
b
b
b
293
298
303
308
313
293
298
303
308
313
0.44
0.66
1.03
1.60
3.55
0.85
1.43
2.25
3.83
5.10
0.57
0.75
1.30
1.85
4.28
1.12
1.89
2.87
4.56
6.06
0.64
0.87
1.51
2.34
4.86
1.25
2.08
3.34
5.24
6.86
0.70
1.02
1.74
2.86
5.22
1.48
2.43
3.89
6.12
7.69
0.78
1.13
1.96
3.20
5.88
1.62
2.75
4.24
6.61
8.69
1.10
1.10
1.10
1.10
1.10
a Constant ionic strength was achieved by using NaHSO4 as an added electrolyte.
b Initial ionic strength is 0.16 mol dmϪ3, but changes in the course of the reaction.

TARTARIC ACID OXIDATION
59
Table VII Activation Parameters for the Tartaric Acid Oxidation by Vanadium(V)
Calculated from the Values of the Apparent Second-Order Rate Constant, kЈ (molϪ1 dm3 sϪ
1
)
obtained from Table VI
Ϫ3
Ϫ
1
Þ
Ϫ
1
Þ
Ϫ
Ϫ1
I/mol cm
Ea/kJ mol
⌬H /kJ mol
⌬S /J mol 1 K
a
77.1
59.3
74.6
56.8
Ϫ57
Ϫ110
1.10
a Initial ionic strength is 0.16 mol dmϪ3, but changes in the course of the reaction.
behaves as a strong monobasic acid for the purpose of
hydrogen and/or bisulfate ions.
presented a linear dependence on the plot of log kobs
against the square root of until 2.75 mol dmϪ3 for the
ionic strength [10,12]. This means that the primary salt
effect [15] is surprisingly applicable to the ionic
The oxidation of tartaric acid, HTA, by vana-
dium(V), V(V), takes place more rapidly in compar-
ison with lactic acid [11] and malic acid [10] indicat-
ing that electron withdrawing substituents at the
␣-position increased the oxidation rate. This was ex-
perimentally verified through the use of a smaller con-
centration of sulfuric acid in the kinetic measurements.
In the previous article [11] we assumed that the
3
strength upper 0.05 mol dmϪ . The reasons that lead
to this behavior are unknown yet. In this reaction the
ionic charge product, zϩzϪ, of 0.56 was obtained by
the variation of the ionic strength from 0.16 to 2.00
mol dmϪ3 (Table IV). From this value alone, it is dif-
ficult to suggest that the vanadium(V) reacts with tar-
taric acid either as the protonated form or in the mo-
lecular form. Furthermore, the plots of ln kobs against
the reciprocal of the dielectric constant of methanol-
water composition [17], with no control or constant
ionic strength, give straight lines. This suggests an ion-
dipole interaction under the assumption that HTA be-
haves as a dipole in methanol, whereas positive ionic
species of V(V) may be considered active in the re-
action. The less polar solvents have a greater loss in
freedom in becoming frozen to the ions than the more
polar solvents, hence the lower rates of reaction in less
polar solvents [19]. The increase in the reaction rate
of the mixture with smaller proportion of methanol
indicates that the activated complex radial term in the
Born expression [20] is less than the sum of the rea-
gents radii in the solution with low dielectric constant
and, thus favors the production of species with higher
dipole moments.
ϩ
pervanadyl ion, VO2 , obtained in acidic solution
3
([Hϩ] Ͼ 0.005 mol dmϪ ) exists in the hydrated form,
ϩ
V(OH)4 [9,18]. In more concentrated solution, the
protonated form, V(OH)3 ϩ is more reactive. The ob-
2
served increase in the reaction rate with the V(V) con-
2
2
centration varying from 0.50 ϫ 10Ϫ to 2.50 ϫ 10Ϫ
mol dmϪ , regardless of the ionic strength in the sul-
furic acid medium (Table II), and with the addition of
sodium bisulfate, varying the ionic strength from 0.16
3
to 2.00 mol dmϪ by approximately four-fold (Table
3
IV), can be attributed to the formation of different spe-
cies produced by the interaction between hydrated va-
nadium (V) and the bisulfate group before the reaction
with tartaric acid, according to the following scheme
K1
2
ϩ
ϩ
V(OH)3 ϩ HSO4Ϫ E
RF V(OH)3HSO4
(1)
The negative values of the entropy of activation
(Table VII) are indicative of the formation of more
rigid activated complexes than the reagents, due to the
approach of like charges which undergo electrostric-
tion with the solvent molecules [20]. A more negative
for the reactions with no control of ionic strength,
where K1 ϭ k1/k1Ј and
K1
RF
Ј
2
ϩ
V(OH)3 ϩ Hϩ ϩ 2HSO4Ϫ E
ϩ
Ϫ
1
mol 1 KϪ
)
V(OH)2(HSO4)2 ϩ H2O (1Ј)
value (Ϫ110 J
and therefore a lower fre-
quency factor (3.0 ϫ 108
dm3
was found
1
1
sϪ
molϪ
)
for fixed ionic strength conditions. This may suggest
that the number of HSO4Ϫ groups in the vanadium(V)
species increases the activated complex stability by
comparison with a less negative value obtained in the
under controlled ionic strength conditions. On the
other hand, when [Hϩ] was varied from 0.05 to 1.00
3
mol dmϪ (Table III), the variation range of the rate
constants with no control of ionic strength was greater
in comparison with those at constant ionic strength and
both showed the trend to a limit value. This yields an
order of 0.60 with respect to [Hϩ] for the former, and
0.28 at constant ionic strength.
1
medium without ionic strength control (Ϫ57 J molϪ
Ϫ
KϪ1), where the oxidant has only one HSO4 group,
leading to the higher activation energy (77.1 kJ
molϪ1). Furthermore, it is possible to suggest that the
oxidation is controlled by entropy at constant ionic
This oxidation and the earlier ones [10–12] have

60
TAKASHIMA, ZIGLIO, AND RONCONI
strength, and by enthalpy (74.6 kJ mol KϪ1) without
ionic control. In addition to these facts, the fractional-
orders with respect to [HTA] and [Hϩ] suggest an
equilibrium for the oxidation with no control of ionic
strength according to
suming the slow step as the rate-determining one, k3 ,
and k3 , respectively, these mechanisms lead to the rate
Ј
law (5) for the oxidation of tartaric acid by V(V) in
sulfuric acid medium with no control of ionic strength,
d[V(V)]t k3K2[Hϩ][HTA][V(V)]t
Ϫ
ϭ
(5)
ϩ
V(OH)3HSO4 ϩ Hϩ
dt
1 ϩ K2[Hϩ][HTA]
K2
R
ϩ
ϩ HOOCCH(OH)CH(OH)COOH E
and,
V(OH)2(HSO4)2 ϩ Hϩ
F (X*2
)
(2)
and
d[V(V)]t k3 K2 [Hϩ][HTA][V(V)]t
Ј
Ј
Ϫ
ϭ
(5Ј)
dt
1 ϩ K2 [Hϩ][HTA]
ϩ
Ј
K2
R
Ј
ϩ
ϩ HOOCCH(OH)CH(OH)COOH E
F (Y*2
)
at constant ionic strength, where the total or analytical
concentrations of V(V) are given by
(2Ј)
with control of ionic strength. It has been shown ear-
lier that most reactions involving V(V) proceed via a
free-radical mechanism [18]. In this investigation the
addition of acrylamide to the reaction mixture has also
been shown to yield the formation of a polymeric
ϩ
ϩ
[V(V)]t ϭ [(X*2 )] ϩ [V(OH)3HSO4
]
or
ϩ
[V(V)]t ϭ [(Y*2 )] ϩ [V(OH)2(HSO4ϩ)2]
product indicating that V(V) behaves as an one-equiv-
ϩ
alent oxidant. Thus, these activated species, (X*2
)
The rate laws, (5) and (5Ј), are in agreement with the
experimental results in sulfuric acid medium, with no
control or under fixed ionic strength, because the in-
vestigated reactions are of first-order with respect to
the vanadium(V) concentration, and of fractional-or-
ders to both tartaric acid and hydrogen ion concentra-
tions. From the slopes and intercepts of the double
reciprocal plots at constant vanadium(V) and sulfuric
acid concentrations (Fig. 2, r Ն 0.9736 and s Յ
and (Y*2ϩ), are decomposed to give the first CO2
through a C9C fission producing a free radical,
HOOCCH(OH)CH(OH)D, and V(IV) in the rate de-
termining step, k3 and kЈ3, respectively, for the oxida-
tion with no control and with constant ionic strength.
H
k3
9
(X*2ϩ) 9
: HOOCCH(OH)CD ϩ CO2
OH
ϩ
ϩ DV(OH)2HSO4 ϩ Hϩ ϩ H2O (3)
and
H
: HOOCCH(OH)CD ϩ CO2
OH
k3
9
(Y*2ϩ) 9Ј
ϩ
ϩDV(OH)(HSO4)2 ϩ Hϩ ϩ H2O (3Ј)
This free radical is rapidly oxidized through another
mol of V(V) and produces one more carbon dioxide,
two moles of formaldehyde and V(IV), according to
the following scheme in the absence of ionic strength
control,
H
k4
HOOCCH(OH)CD ϩ V(OH)3HSO4ϩ 9
9: 2HCOH
Figure 2 A double reciprocal plot of the pseudo-first-order
rate constant, kobs, and tartaric acid concentration, [HTA] for
OH
ϩ
Ϫ2
Ϫ3
ϩ CO2 ϩ DV(OH)2HSO4 ϩ H2O (4)
[V(V)] ϭ 1.00 ϫ 10 mol dm and [H2SO4] ϭ 0.15 mol
Ϫ
dm 3 at 303 K. (᭿) No control of the ionic strength and (᭹)
Ϫ
and, in a similar way at constant ionic strength. As-
constant ionic strength (I ϭ 1.10 mol dm 3).

TARTARIC ACID OXIDATION
61
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18.8089) the equilibrium constant, K2 was determined
as ca. 6.0 molϪ2 dm6 at 303 K, while the rate constant
for rate-determining constant, k3 , was obtained as
1
1
3.38 ϫ 10Ϫ3 molϪ dm3 sϪ with no control of ionic
strength control. When this was adjusted to 1.10 mol
3
2
dmϪ , these values become 25.1 molϪ dm6 and
4.51 ϫ 10Ϫ3 molϪ dm3 sϪ , respectively. In conclu-
sion, these results and the experimental rate laws show
that the oxidation mechanism of tartaric acid involves
different active species of vanadium(V) in sulfuric
acid medium and depend on the ionic strength condi-
tions of the experiments.
1
1
11. R. M. Clementin and K. Takashima, Qu´ımica Nova, 16,
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Cliffs, 1989, p. 164–168.
This work was possible through the financial support from
PADCT/FINEP, PADCT/CAPES, CNPq, and CPG/UEL;
from CNPq through the Scientific Initiation Scholarship
(C.M.Z. and C.M.R.) and a Research Fellowship (K.T.). We
thank Prof. Jose´ Manuel Riveros (University of Sa˜o Paulo)
for valuable suggestions concerning this study.
16. G. G. Manov, R. G. Bates, W. J. Hamer, and S. F. Acree,
J. Am. Chem. Soc., 65, 1765 (1943).
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