Kinetics and mechanism of the oxidation of some carboxylates by a nickel(III) oxime-imine complex

Article abstract of DOI:10.1002/(SICI)1097-4601(1997)29:3<225::AID-KIN10>3.0.CO;2-9

The kinetics of the oxidation of formate, oxalate, and malonate by [Ni^III(L¹)]²⁺ (where HL¹ = 15-amino-3-methyl-4,7,10,13-tetraazapentadec-3-en-2-one oxime) were carried out over the regions pH 3.0-5.75, 2.80-5.50, and 2.50-7.58, respectively, at constant ionic strength and temperature 40°C. All the reactions are overall second-order with first-order on both the oxidant and reductant. A general rate law is given as - d/dt[Ni^III(L¹)²⁺] = k_obs[Ni^III(L¹)²⁺] = (k_d + nk_s[R])[Ni^III(L¹)²⁺], where k_d is the auto-decomposition rate constant of the complex, k_s is the electron transfer rate constant, n is the stoichiometric factor, and R is either formate, oxalate, or malonate. The reactivity of all the reacting species of the reductants in solution were evaluated choosing suitable pH regions. The reactivity orders are: k_HCOOH > k_HCOO(-); k_H(2)ox > k_Hox(-) > k_ox(2-), and k_H(2)mal > k_Hmal(-) < k_mal(2-) for the oxidation of formate, oxalate, and malonate, respectively, and these trends were explained considering the effect of hydrogen bonded adduct formation and thermodynamic potential.

Full text of DOI:10.1002/(SICI)1097-4601(1997)29:3<225::AID-KIN10>3.0.CO;2-9

Kinetics and Mechanism of
the Oxidation of Some
Carboxylates by a
Nickel(III) Oxime-Imine
Complex
BASUDEB SAHA, AMITAVA DUTTA, SUMANA GANGOPADHYAY, PRADYOT BANERJEE
Department of Inorganic Chemistry, Indian Association for the Cultivation of Science, Calcutta 700 032, India
Received 9 May 1996; accepted 29 July 1996
III
1
2
ϩ
ABSTRACT: The kinetics of the oxidation of formate, oxalate, and malonate by [Ni (L )]
(where HL1 ϭ 15-amino-3-methyl-4,7,10,13-tetraazapentadec-3-en-2-one oxime) were carried
out over the regions pH 3.0–5.75, 2.80–5.50, and 2.50–7.58, respectively, at constant ionic
strength and temperature 40ЊC. All the reactions are overall second-order with first-order on
both the oxidant and reductant. A general rate law is given as Ϫd/dt[NiIII(L1)2ϩ] ϭ
k
obs[NiIII(L1)2ϩ] ϭ (kd ϩ nks[R])[NiIII(L1)2ϩ], where kd is the auto-decomposition rate constant
of the complex, ks is the electron transfer rate constant, n is the stoichiometric factor, and R is
either formate, oxalate, or malonate. The reactivity of all the reacting species of the reduc-
tants in solution were evaluated choosing suitable pH regions. The reactivity orders are:
k
HCOOH Ͼ kHCOOϪ; kH ox Ͼ kHoxϪ Ͼ kox2Ϫ, and kH mal Ͼ kHmalϪ Ͻ kmal2 for the oxidation of for-
Ϫ
2
2
mate, oxalate, and malonate, respectively, and these trends were explained considering the
effect of hydrogen bonded adduct formation and thermodynamic potential. © 1996 John Wiley
& Sons, Inc. Int J Chem Kinet 29: 225–230, 1997.
redox studies including organic and pure inorganic
reductants have been made involving this ion [2–7],
a broad knowledge concerning the electron transfer in
simple carboxylates is yet to be explored. Among the
simple carboxylates, formate is of significant interest
because of its relevance to plant metabolism leading
to oxidative decarboxylation [8,9]. The importance of
oxalate is due to its use as a quencher in solar photo-
chemistry [10–13] and as a potential electron donor
in other excited state electron transfer reactions. On
the other hand, the reactions of malonic acid with a
variety of oxidants also showed versatile mechanistic
features [14–18]. The present investigation concerns
INTRODUCTION
The synthesis and characterization of oxime-imine
ligand like HL1(I) (15-amino-3-methyl-4,7,10,13-
tetraazapentadec-3-ene-2-one oxime) which can sta-
bilize the oxidation state of nickel upto ϩ 3 is well
documented in the literature [1]. This [Ni (L )]
complex shows interesting chemical features in the
redox transformations [1,2]. Although a number of
2
III
1
ϩ
Correspondence to: P. Banerjee
Contract grant Sponsor: Council of Scientific and Industrial Re-
search
2
III
1
ϩ
the study of redox reactions of [Ni (L )] complex
with formate, oxalate, and malonate.
© 1997 John Wiley & Sons, Inc.
CCC 0538-8066/97/030225-06

226
SAHA ET AL.
III
1
2
ϩ
[Ni (L )] complex was not possible due to slow-
ness of the reactions and time dependent decomposi-
tion of the complex. Their product analyses were car-
ried out very carefully in order to throw some light on
the stoichiometries. For the oxidation of formate and
oxalate the product was carbon dioxide as tested by
the usual standard method. Analysis of the oxidation
product of malonic acid was performed by reacting
NOH
H2N
H
N
H
N
N
NH
(I)
III
1
2
ϩ
equimolar concentrations of [Ni (L )]
complex
EXPERIMENTAL
Materials and Reagents
1
and the reducing agent (0.05 mol LϪ ) as reported
earlier [18]. The product was condensed for an hour
[19] with an equimolar quantity of o-phenylenedi-
amine in a minimum volume of dry ethanol. The
solid residue was obtained on concentrating and cool-
ing overnight in a deep freeze. It was recrystallized
from anhydrous ethanol and dried in a desiccator. The
melting point of the compound determined was
265.8ЊC which is identical to that of 2-hydroxy-3-iso-
quinoxaline carboxylic acid [20] obtained from the
condensation of mesoxalic acid with o-phenylenedi-
amine. The product also showed a positive 2,4-dini-
trophenyl hydrazine test. Mesoxalic acid has been
identified as a product in many of the oxidation reac-
tions of malonic acid where a 4:1 stoichiometry with
respect to the oxidant prevails [21,22]. The stoichio-
metric equations for the oxidation of formate, ox-
alate, and malonate at pH 4.5 can be represented as:
(15-amino-3-methyl-4,7,10,13-tetraazapentadec-3-
ene-2-one oximato) nickel(III) perchlorate, [NiIII
(L1)](ClO4)2 was prepared by the method reported by
Chakravorty et al. [1]. Analytical data were satis-
factory (found: C, 27.30; H, 5.10; N, 15.85%, re-
quired for C12H27Cl2N6NiO9 :C, 27.35; H, 5.30; N,
15.80%, respectively). The complex solutions were
standardized spectrophotometrically using the litera-
ture absorption coefficient: ␭/nm(⑀/L molϪ cmϪ )
498[(2.7 Ϯ 0.20) ϫ 103]. Sodium formate (E.
Merck), sodium malonate (E. Merck) and oxalic acid
were of reagent grade and recrystallized twice before
use. Doubly recrystallized sodium perchlorate (Fluka
AG) was used to maintain the ionic strength of the
medium. Doubly recrystallized sodium acetate and
sodium dihydrogen phosphate, and acetic acid, per-
chloric acid, and sodium hydroxide, all of reagent
grade, were used to maintain the pH of the medium.
All the solutions were prepared freshly prior to ki-
netic study using double distilled water.
1
1
2[NiIII(L1)]2ϩ ϩ HCOOH
2[NiIII(L1)]2ϩ ϩ H2C2O4
2[NiII(HL1)]2ϩ ϩ CO2
2[NiII(HL1)]2ϩ ϩ 2CO2
COOH
4[NiIII(L1)]2ϩ ϩ H2C
ϩ H2O
COOH
COOH
Kinetic Measurements
4[NiII(HL1)]2ϩ ϩ O
C
Kinetics of the reactions were followed using a uv-vis
spectrophotometer (uv-2100, Shimadzu, Japan)
equipped with thermostated cell compartments by
monitoring the decrease in optical density at 498 nm.
pH measurements were carried out with a Systronics
(model 335, India) digital pH-meter. Although mea-
sured pH is usually defined in terms of hydrogen ion
activity, the hydrogen ion concentration for each so-
lution was obtained by calibrating the pH electrode
with analytically prepared solutions maintained to the
desired ionic strength. The reaction temperature
(Ϯ 0.1ЊC) was controlled by circulating water around
the reservoir for the reaction solutions from a Haake
F3 thermostat.
COOH
Polymerization Study
All the three reactions were tested for the generation
of free radicals during the course of the reaction by
the method reported earlier [23]. The precipitation of
a white polymer of acrylonitrile indicates that all the
reactions proceed through the generation of free radi-
cals.
RESULTS AND DISCUSSION
The kinetics of the oxidation of formate, oxalate, and
Stoichiometry and Reaction Products
III
1
2
ϩ
malonate by [Ni (L )] were carried out under
pseudo-first-order conditions with excess of reductant
(Ͼ 10 times) over the complex at a constant tempera-
The direct determination of the stoichiometry for
the oxidation of formate, oxalate, and malonate with

OXIDATION OF SOME CARBOXYLATES
227
Table I Pseudo-First-Order Rate Constants for the
Reduction of [NiIII(L1)]2ϩ by Formate, Oxalate, and
Malonate at Different Reductant Concentrations with
[NiIII(L1)2ϩ] ϭ 1.0 ϫ 10Ϫ mol LϪ and [OAcϪ] ϭ 0.02
4
1
mol LϪ1 at 40ЊC
1
1
Reductant
Formatea
[Reductant], mol LϪ
104kobs, , sϪ
0.10
0.30
0.50
0.70
1.00
0.38
0.48
0.64
0.74
0.91
Figure 1 Variation of ks as a function of Ϫlog[Hϩ] for
0.001
0.003
0.005
0.007
0.010
4.60
10.20
14.80
20.10
28.30
the reduction of [NiIII(L1)]2ϩ by (a) formate and (b) oxalate
at 40ЊC with [NiIII(L1)2ϩ] ϭ 1.0 ϫ 10Ϫ mol LϪ , I ϭ 1.0
4
1
Oxalateb
1
1
mol LϪ for formate and 0.20 mol LϪ for oxalate, and
1
[buffer] ϭ 0.02 mol LϪ . The solid line represents calcu-
lated ks values and the experimental values shown by
points.
0.10
0.30
0.50
0.70
0.90
0.46
0.86
0.92
1.35
1.65
Malonatec
a I ϭ 1.00 mol LϪ1 and kd ϭ (3.50 Ϯ 0.50) ϫ 10Ϫ5 sϪ1 at pH
b I ϭ 0.20 mol LϪ1 and kd ϭ (2.80 Ϯ 0.20) ϫ 10Ϫ4 sϪ1 at pH
c I ϭ 1.00 mol LϪ1 and kd ϭ (4.00 Ϯ 0.20) ϫ 10Ϫ5 sϪ1 at pH
5.0.
3.0.
5.0.
Oxidation of Formate
III
1
2
ϩ
Oxidation of formate by [Ni (L )] was carried
III
1 2
ϩ
out in the range pH 3.0–5.75 with [Ni (L ) ] ϭ
4
1
1
1.0 ϫ 10Ϫ mol LϪ , [formate] ϭ 0.1–1.0 mol LϪ ,
1
and I ϭ 1.0 mol LϪ (NaClO4).
ks(ϭ kobs Ϫ kd / 2[formate]) vs. Ϫlog[Hϩ] is a curve
A plot of
of decreasing slope (Fig. 1(a)). The complex
ture 40ЊC. The plots of Ϫlog At vs. t (where At repre-
sents the absorbance at time t) are linear for more
than 75% of the total reactions. The pseudo-first-
order rate constants (kobs) were evaluated from the
slopes of the straight lines and vary within the error
limit of Ϯ 3%. Again the plots of kobs at different re-
ductant concentrations yield straight lines with a
nearly common intercept on the rate axis at a certain
pH which may indicate the auto-decomposition of the
complex at that particular pH. Corresponding data are
displayed in Table I. A general rate law consistent to
the oxidation of all these three carboxylates may be
given as
III
1
2
ϩ
[Ni (L )] is a strong acid (pKm Ͻ 1.0) [2], the re-
acting species of the complex in our experimental pH
region would be the deprotonated one. Electrochemi-
cal studies showed a quasi-reversible wave with E0 ϭ
0.91 V (vs. NHE) and given as
[NiIII(L1)]2ϩ ϩ e ϩ Hϩ EF [NiII(HL1)]2ϩ (2)
II
1
2
ϩ
The pKm of the [Ni (HL )] complex is represented
as
Km
[NiII(HL1)]2ϩ EF [NiII(L1)]ϩ ϩ Hϩ; pKm ϭ 6.8
(3)
d
Ϫ
[NiIII(L1)2ϩ] ϭ kobs[NiIII(L1)2ϩ
]
The potential of the related couple (4) has been
estimated to be 0.49 V by using the value of pKm
(eq. (3))
dt
ϭ (kd ϩ nks[R])[NiIII(L1)2ϩ] (1)
[NiIII(L1)]2ϩ ϩ e EF [NiII(L1)]ϩ
(4)
where kd is the auto-decomposition rate constant of
the complex, found to be pH-dependent and studied
separately earlier [7], ks is the electron transfer rate
constant, R is either formate, oxalate or malonate, n
is the stoichiometric factor which is 2 for both for-
mate and oxalate oxidations, and 4 for malonate oxi-
dation.
Thus, the proton dependence of rate in the experi-
mental range of pH can be unambiguously ascribed
to the protic equilibrium of formic acid [24],
K1
HCOOH EF HCOOϪ ϩ Hϩ; pK1 ϭ 3.53 (5)

228
SAHA ET AL.
As the reaction rates (k0) fall with Ϫlog[Hϩ], it is ex-
pected that HCOOH is more reactive compared to the
deprotonated species, HCOOϪ. This is contrary to
our previous observations for formate oxidations
[18,25]. Assuming [Ni (L )] as the sole reacting
species, a suitable reaction scheme compatible to the
experimental findings could be framed as
K1(ϪHϩ
)
K2(ϪHϩ
)
COOH
COOH
(H2ox)
COOϪ
COOH
(HoxϪ)
COOϪ
COO
(11)
(ox2Ϫ
)
III
1
2
ϩ
with pK1 ϭ 1.20 and pK2 ϭ 4.21. The reactive
species of oxalate are likely to be H2ox, HoxϪ, and
ox2Ϫ in the experimental pH range and a plausible re-
action scheme could be framed as:
k6
[NiIII(L1)]2ϩ ϩ HCOOH 9: [NiII(L1)]ϩ
ϩ
и
ϩ HCOO ϩ H (6)
k12
[NiIII(L1)]2ϩ ϩ H2ox 9: [NiII(L1)]ϩ
k 7
[NiIII(L1)]2ϩ ϩ HCOOϪ 9: [NiII(L1)]ϩ
ϩ
и
ϩ Hox ϩ H
(12)
(13)
(14)
и
ϩ HCOO
(7)
(8)
k13
[NiIII(L1)]2ϩ ϩ HoxϪ 9: [NiII(L1)]ϩ ϩ Hox
и
fast
[NiIII(L1)]2ϩ ϩ radical 9: [NiII(L1)]ϩ
ϩ CO2
k14
[NiIII(L1)]2ϩ ϩ ox2Ϫ 9: [NiII(L1)]ϩ ϩ ox
и
and corresponding rate law could be derived as
The rate of the free radicals would be the same as
mentioned earlier and the rate law can be derived as:
k6[Hϩ] ϩ k7K1
ks ϭ
(9)
[Hϩ] ϩ K1
k12[Hϩ]2 ϩ k13 K1[Hϩ] ϩ k14 K1 K2
Since pKm of the complex [NiII(HL1)]2 is 6.8, the
reduction product of the complex would be the proto-
nated species, in the experimental pH range.
ϩ
ks ϭ
(15)
[Hϩ]2 ϩ K1[Hϩ] ϩ K1 K2
Equation (15) has been solved by means of a simplex
optimization program and the evaluated parameters
[NiII(L1)]ϩ ϩ Hϩ
9
9
rap:id [NiII(HL1)]2ϩ
(10)
1
1
1
are: k12 ϭ (6.00 Ϯ 0.18) ϫ 10Ϫ L molϪ sϪ , k13 ϭ
3
1
1
(2.27 Ϯ 0.07) ϫ 10Ϫ L molϪ sϪ , k14 ϭ (2.21 Ϯ
In the absence of any monomer, the free radicals
would react in a fast step with another [NiIII(L1)]2ϩ to
yield CO2 and Ni(II) products (eq. 8). Equation (9)
was solved by means of a nonlinear least-squares pro-
gram and the evaluated parameters are:
4
1
1
0.09) ϫ 10Ϫ L molϪ sϪ , K1 ϭ (4.25 Ϯ 0.15) ϫ
3
1
10Ϫ mol LϪ (pK1 ϭ 2.37), and K2 ϭ (4.35 Ϯ
0.15) ϫ 10Ϫ mol LϪ1 (pK2 ϭ 4.36).
5
Oxidation of Malonate
4
1
1
k6 ϭ (6.75 Ϯ 0.04) ϫ 10Ϫ L molϪ sϪ , k7
5
1
1
ϭ (1.01 Ϯ 0.14) ϫ 10Ϫ L molϪ sϪ , and K1
The kinetics of the oxidation of malonate was studied
using variable concentrations of malonate (0.10–0.90
4
ϭ (5.35 Ϯ 0.06) ϫ 10Ϫ mol LϪ1 (pK1 ϭ 3.27)
1
mol LϪ ) in the range pH 2.50–7.58 using fixed con-
4
1
centration of the complex (1.0 ϫ 10Ϫ mol LϪ ) at a
It is seen that there is an excellent agreement between
the experimental and calculated ks values (Fig. 1(a))
at each [Hϩ] but there is a shift in pK1 value for
HCOOH towards the lower side.
1
constant ionic strength [1.0 mol LϪ (NaClO4)]. A
plot of Ϫlog[Hϩ] vs. ks (Fig. 2) showed a decrease in
rate with Ϫlog[Hϩ], attains a minimum at around pH
5.0–5.75 (Ϫlog[Hϩ] ϭ 4.89–5.61) and then rate in-
creases steadily with Ϫlog[Hϩ]. The pK’s of malonic
acid are, pK1 ϭ 2.88 and pK2 ϭ 5.68 [26] owing to
the equilibria,
Oxidation of Oxalate
III
1
2
ϩ
Oxidation of oxalic acid by [Ni (L )] was carried
K1
K2
COOH
COOH
COOϪ
COOϪ
COOϪ
out in the range pH 2.80–5.50 at [ox] ϭ (1.0–10.0)
ϪHϩ
ϪHϩ
3
1
4
ϫ 10Ϫ mol LϪ , [NiIII(L1)]2ϩ ϭ 1.0 ϫ 10Ϫ mol
H2C
H2C
H2C
ϩHϩ
ϩHϩ
1
1
LϪ , and I ϭ 0.2 mol LϪ (NaClO4). A plot of
Ϫlog[Hϩ] vs. ks is also a curve of decreasing slope
(Fig. 1(b)]. The pK’s [26] of oxalic acid are desig-
nated as,
COOH
(H2mal)
(HmalϪ)
(mal2Ϫ
)
(16)

OXIDATION OF SOME CARBOXYLATES
229
Equation (21) was solved by a nonlinear least-squares
computer-fit program and the evaluated values are as:
4
1
1
k
17 ϭ (5.06 Ϯ 0.04) ϫ 10Ϫ L molϪ sϪ , k18
5
1
1
ϭ (1.75 Ϯ 0.10) ϫ 10Ϫ L molϪ sϪ , and K1
4
ϭ (9.80 Ϯ 0.15) ϫ 10Ϫ mol LϪ1 (pK1 ϭ 3.0).
Similar approach was made in the pH range
5.80–7.58 to evaluate the rate parameters (k18 , k19)
and K2 and the rate law takes the form
k18[Hϩ] ϩ k19 K2
ks ϭ
(22)
[Hϩ] ϩ K2
Figure 2 Variation of ks as a function of Ϫlog[Hϩ] for the
oxidation of malonate at 40ЊC with [NiIII(L1)2ϩ] ϭ 1.0 ϫ
The evaluated parameters are: k18 ϭ (1.67 Ϯ 0.28) ϫ
4
1
1
10Ϫ mol LϪ , I ϭ 1.0 mol LϪ , and [buffer] ϭ 0.02 mol
5
1
1
4
10Ϫ L molϪ sϪ , k19 ϭ (1.05 Ϯ 0.1) ϫ 10Ϫ
L
1
LϪ . The curve has been drawn through calculated ks val-
1
1
7
1
molϪ sϪ , and K2 ϭ (3.85 Ϯ 0.12) ϫ 10Ϫ mol LϪ
(pK2 ϭ 6.41). All these evaluated rate and equilib-
rium constants were incorporated in the general rate
expression (eq. (20)) and the values of ks were calcu-
lated at each pH. These match excellently with the
experimentally determined values (Fig. 2).
ues, while circles denote experimental points.
In the range pH 2.50–7.58 the reacting species of
malonate are expected to be H2mal, HmalϪ, and
Nickel(III) ion in [NiIII(L1)]2ϩ is in a pseudo-octa-
hedral geometry (t2g6eg ) with the odd electron in dz2
1
2
Ϫ
mal , and their reactions towards the complex
orbital. This complex is assumed to have a very low
substitution lability though there is no accurate esti-
mate for it. Due to low substitution lability, it is ex-
pected that the electron transfer reactions of
III
1
2
ϩ
[Ni (L )] , are portrayed as,
k17
[NiIII(L1)]2ϩ ϩ H2mal 9: [NiII(L1)]ϩ
ϩ
и
[NiIII(L1)]2 would occur via outer-sphere pathway.
ϩ H
ϩ
ϩ Hmal
(17)
k18
III
1
2
ϩ
However, in most of the oxidations by [Ni (L )] ,
e.g., of ascorbic acid [6], benzene diols [2], S(IV)
and Se(IV) [4], the reactions were proposed to take
place via the formation of hydrogen-bonded interme-
diate which is supported by the higher reactivity of
protonated species over the less protonated or depro-
tonated one. This is particularly apparent for S(IV)
and Se(IV) oxidations by [NiIII(L1)]2ϩ [4], where the
[NiIII(L1)]2ϩ ϩ HmalϪ 9: [NiII(L1)]ϩ ϩ Hmal
и
(18)
k19
[NiIII(L1)]2ϩ ϩ mal2Ϫ 9: [NiII(L1)]ϩ ϩ mal
и
(19)
The radicals thus produced react with another NiIII
complex to give NiII complex and mesoxalic acid as
the final products. If it is assumed that the reaction
stoichiometry (4:1) is valid under the reaction condi-
tion, the corresponding rate law could be written as,
overall reactivity order was found to be SO x H O
и
2
2
2
Ϫ
Ϫ
Ͼ HSO3Ϫ Ͻ SO3 and H2SeO3 Ͼ HSeO3 , respec-
tively, and the formation constant values follow the
2
trend: SO x H O Ͼ HSO Ϫ Ͼ SO3 Ϫ. In the pre-
и
2
2
3
k
obs Ϫ kd
k17[Hϩ]2 ϩ k18K1[Hϩ] ϩ k19K1K2
sent study too, the oxidation rates were found to de-
crease with increase in pH (except for the malonate
oxidation which is comparable to S(IV) oxidation)
resulting in the higher reactivity of the protonated
species over the less protonated and completely de-
protonated ones. In fact the reactivity orders are:
ks ϭ
ϭ
4[mal]
[Hϩ]2 ϩ K1[Hϩ] ϩ K1K2
(20)
The kinetic and equilibrium parameters of eq. (20)
could be evaluated by considering suitable pH ranges.
Thus in the range pH 2.50–5.50, the reactions (17)
and (18) are likely to occur and the rate law reduces
to
k
HCOOH Ͼ kHCOOϪ ; kH ox Ͼ kHoxϪ Ͼ k Ϫ , and k
ox H mal
2 2
2
Ͼ kHmalϪ Ͻ k Ϫ . The order of reactivity of carboxy-
2
mal
lates is contrary to the previous results [25,27,18]
where the reverse trend was noted and an explanation
was furnished based on thermodynamic potential
considerations. In the present study, the effect of
k17[Hϩ] ϩ k18 K1
ks ϭ
(21)
[Hϩ] ϩ K1

230
SAHA ET AL.
7. A. Dutta, B. Saha, M. Ali, and P. Banerjee, Int. J.
Chem. Kinet., communicated.
8. J. Halliwell, Biochem. J., 138, 177 (1974).
9. E. R. Still and P. Wikberg, Inorg. Chim. Acta, 46, 153
(1980).
H-bonded adduct formation seems to predominate
over the thermodynamic potential considerations pro-
viding a lower energy pathway for electron transfer
from the reductant to [Ni (L )] . The formation of
H-bonded adduct is expected to occur between the
III
1
2
ϩ
10. K. Zahir, J. H. Espenson, and A. Bakac, Inorg. Chem.,
27, 3144 (1988).
carboxylate hydrogen and oximato oxygen atom
l
("N9OϪ) of the coordinating ligand. The esti-
mates for the formation constant of the H-bonded
adducts, however, is not possible as the reactions do
not proceed through rate saturation in our experimen-
tal reductant concentration but a definite impact on
the pK values of the reductants is noted in some
cases.
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Of interest is the comparison of the reactivity of
2
mal to those of HmalϪ and H2mal leading one to
believe that thermodynamic potential values may ex-
ert well predominance here over the hydrogen-bond-
ing effect. Similar behavior to this was observed in
the study of the S(IV) oxidation with [NiIII(L1)]2ϩ [4]
Ϫ
2
Ϫ
where SO3 shows a higher reactivity compared to
So2 x H2O and HSO3Ϫ species.
и
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P. Banerjee thanks the Council of Scientific and Industrial
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for a fellowship.
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tional Conference on Mechanism of Reactions in Solu-
tion, The University of Kent at Canterbury, United
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