Manganese(III) oxidation of alanine in aqueous ethanoic acid medium: A kinetic and mechanistic study

Article abstract of DOI:10.1080/15533174.2011.609214

Astock solution of manganese(III) acetate in aqueous acetic acid has been prepared using a standard electrochemical method. The nature of the oxidizing species present in manganese(III) solutions has been characterized by spectrophotometric and redox potential measurements. Kinetics of oxidation of DL-alanine (Ala) by manganese( III) in aqueous acetic acid solution at 313Khave been investigated. The experimental rate law is: rate = k'[Mn(III)] ^0.5[Ala]^0.9 [OAc-]y^0.5/[H+]^0.5, where each species shows a fractional order. Effects of varying ionic strength, solvent composition, and the added ions such as fluoride, chloride, and perchlorate have been investigated. Activation parameters have been evaluated using Arrhenius andEyring plots.Amechanism consistentwith the observed kinetic data has been proposed and discussed. A suitable rate law has been derived based on the mechanism. Copyright Taylor & Francis Group, LLC.

Full text of DOI:10.1080/15533174.2011.609214

Synthesis and Reactivity in Inorganic, Metal-Organic, and Nano-Metal Chemistry, 41:1331–1337, 2011
Copyright ꢀC Taylor & Francis Group, LLC
ISSN: 1553-3174 print / 1553-3182 online
DOI: 10.1080/15533174.2011.609214
Manganese(III) Oxidation of Alanine in Aqueous Ethanoic
Acid Medium: A Kinetic and Mechanistic Study
1
2
3
C. S. Chidan Kumar, S. Chandraju, and Netkal M. Made Gowda
1
Department of Chemistry, G. Madegowda Institute of Technology, Bharathi Nagara, Mandya,
Karnataka, India
Department of Studies in Sugar Technology, Sir M. Visweswaraya Post Graduate Center,
2
University of Mysore, Tubinakere, Mandya, Karnataka, India
3
Department of Chemistry, Western Illinois University, One University Circle, Macomb, Illinois, USA
be ingested. Colored manganese compounds are used industri-
A stock solution of manganese(III) acetate in aqueous acetic acid ally as pigments and at higher oxidation states the metal ions are
has been prepared using a standard electrochemical method. The used as oxidants.[
1–4]
Manganese(II) ions function as cofactors
nature of the oxidizing species present in manganese(III) solutions
for a number of enzymes and the element is a required trace
mineral for all living organisms. There is a pronounced impor-
tance of manganese in growth, bone development, reproduction,
has been characterized by spectrophotometric and redox potential
measurements. Kinetics of oxidation of DL-alanine (Ala) by man-
ganese(III) in aqueous acetic acid solution at 313 K have been inves-
ꢀ
0.5
0.9
tigated. The experimental rate law is: rate = k [Mn(III)] [Ala]
mitochondrial function, and central nervous system. It is also
−
y0.5
+ 0.5
[1,2]
[
OAc ] /[H ] , where each species shows a fractional order. Ef- important in the photosynthetic oxygen evolution in plants.
fects of varying ionic strength, solvent composition, and the added
ions such as fluoride, chloride, and perchlorate have been investi-
gated. Activation parameters have been evaluated using Arrhenius
and Eyring plots. A mechanism consistent with the observed kinetic
There are several reports of oxidation of various substrates by
Mn(III) in acid solutions containing such salts as perchlorate,
sulfate, acetate, and pyrophosphate.[
7–13]
A study of the kinetics
data has been proposed and discussed. A suitable rate law has been and mechanism of oxidation of DL-Ala by Mn(III) in aqueous
derived based on the mechanism.
acetic acid solutions is reported here.
Keywords alanine, kinetics, manganese(III), mechanism, oxidation,
EXPERIMENTAL
reduction
Preparation of Manganese(III) Acetate Stock Solution
A stock solution of manganese(III) acetate was prepared us-
INTRODUCTION
Oxidations of biologically significant organic substrates by
manganese(III) are of special importance.
[14]
ing the standard method
of anodic oxidation of a solution
containing 0.06 M manganese(II) acetate and 0.5 M sodium ac-
[
1,2]
These oxidations
etate in aqueous acetic acid (90% v/v; 0.500 L) in an undivided
[
3]
[4,5]
[6,7]
2
include reactions of oximes, vitamins,
and amino acids.
cell. Platinum foil anode (generation area 4.0 cm ) and a thin
2
Amino acids serve important functions in biological systems and
play significant roles in metabolism. They are also employed
in biochemical, microbiological, and nutritional investigations.
Alanine (CH3CH(NH2)COOH, Ala) is an α-amino acid that is
used by the body to build protein structures. Since it is not syn-
thesized in the animal body, Ala or proteins containing it must
platinum spiral cathode (effective generation area < 0.2 cm )
were used to perform electrolysis at the cell voltage of approxi-
−2
mately 6 V and a current density of 2 mA cm . The electrolysis
was continued with stirring until the manganese(III) concentra-
tion rose to approximately 10 mM. The resultant manganese(III)
solution was iodometrically standardized using starch indicator
near the endpoint. Further electrolysis resulted in the precip-
itation of manganese(III) acetate due to its limited solubility.
The presence of sodium acetate not only enhances the conduc-
tivity of the electrolyte solution but also increases the solubil-
Received 29 April 2011; accepted 26 July 2011.
The authors are grateful to Drs. Jin Jin and Shaozhong Zhang,
Department of Chemistry, Western Illinois University, Macomb, IL, ity of manganese(III) acetate by forming the complex species
−
for their helpful suggestions on the reaction mechanism. The authors
are also grateful to the Western Illinois University Research Council
for support.
Address correspondence to Netkal M. Made Gowda, Department
of Chemistry, Western Illinois University, One University Circle, Ma-
comb, IL 61455, USA. E-mail: gn-made@wiu.edu
[
Mn(OAc)4] . Optimal conditions for the use of electrochemi-
cal generation from platinum electrodes include 85–95% HOAc
−2
content (v/v), 1–5 mA cm current density, >0.05 M man-
ganese(II), and 0.5 M sodium acetate. The current efficiency
achieved was 79%. It was observed that a stock solution of
1331

1332
C. S. CHIDAN KUMAR ET AL.
.00 × 10− M Mn(III) acetate in aqueous acetic acid (90%, tain a known acid content), sodium acetate (to maintain the total
2
1
v/v) was stable for 2–3 days and can be used for analytical and ionic strength constant), manganese(II) (to maintain a large mole
kinetic studies. However, freshly prepared solutions were al- excess over Mn(III)), and water (to keep the total volume con-
ways preferred. The solution of manganese(III) should consist stant) were mixed in stoppered boiling tubes. The mixtures were
of a large excess of manganese(II) to suppress the dispropor- thermally equilibrated in a water bath at a constant temperature
tionation of the former, as shown in the following equilibrium: (313 K). To each mixture in the tube, was added an aliquot of
preequilibrated standard solution of manganese(III) acetate to
give a known overall concentration. The progress of the reaction
2
Mn(III)(aq) ꢀꢁ Mn(II)(aq) + Mn(1 V)(aq)
[1]
was monitored for at least two half-lives. Aliquots of the reaction
mixture withdrawn at regular time intervals were iodometrically
titrated for the unreacted manganese(III) concentration in each
sample with a standard sodium thiosulfate solution using starch
A chromatographically pure DL-alanine sample (Sisco Re-
search Laboratories, India) was further assayed for the amino
[
15]
acid content by standard methods. A stock aqueous solution
of Ala (0.400 M) was prepared and used as needed. All other
chemicals used were of analytical grade. Double-distilled water
was employed for preparing aqueous solutions.
ꢁ
indicator near the endpoint. Rate constants (k or kobs), calcu-
1/2
lated from the slopes of linear pseudo-half order plots of [titer]
or [Mn(III] versus time, were reproducible within ± 3% error.
1/2
Reaction Stoichiometry and Product Analysis
Preliminary Studies
Visible spectra of freshly prepared manganese(III) acetate so-
lutions at different concentrations of acetic acid and sodium ac-
etate showed the presence of a complex species, [Mn(OAc)4] .
This species, formed at higher acetate concentrations, probably
resulted from the ligand substitution of acetate ions for water
ligands in the complex ion [Mn(OH2)4] . The evidence was in
the visible spectrum of manganese(III), which showed an en-
hanced absorption at λmax = 470 nm. When Ala was mixed
with the Mn(III) acetate solution, the peak blue shifted to 446.0
nm, suggesting a complexation reaction between Mn(III) and
Ala leading to the formation of major products Mn(II) and alde-
hyde. The formal redox potential (E 0) of the Mn(III)–Mn(II)
couple, which is a measure of the oxidizing power of Mn(III),
generally decreases on complexation.
manganese(III) were anodically generated in a given volume of
aqueous acetic acid and the electrode potentials were measured.
The Nernst equation, where Eo(V vs. SCE) represents the potential
of the dynamic equilibrium between the oxidized and reduced
forms of the metal ions established at the platinum electrode
A known amount of Ala was allowed to react completely
with five- to sixfold mole excess of manganese(III) acetate in
the presence of excess manganese(II) in an aqueous acetic acid
−
(60%, v/v) solution at 313 K for 24 h. The concentration of
the unreacted manganese(III) solution was determined by iodo-
metric titrations. Rssults showed that 2 moles of Mn(III) were
consumed by 1 mole of Ala, giving a stoichiometric ratio of 2:1.
The oxidation product, acetaldehyde, in the reaction mixture
was identified by the melting point (165–168◦C) of its 2,4-
3
+
[
16]
[
15]
dinitrophenylhydrazine(DNP) test
and by Tollen’s reagent
test.[ The conventional lime water test was positive for the
presence of CO2. The reduction product, Mn(II), in the reaction
mixture was observed spectrophotmetrically by the disappear-
ance of the Mn(III) absorption peak at 470 nm in its visible
spectrum.
16]
ꢁ
[
16,17]
Varying amounts of
Test for Free Radicals
The Ala–Mn(III) reaction mixtures were tested for the pres-
ence of free radicals using monomers such as acrylonitrile and
methyl methacrylate (10% solution).[ The addition of the
monomer to the reaction mixture, after it was purged with ni-
trogen, in the dark resulted in the precipitation of a polymer
product. This positive response indicated the in situ generation
of free radicals in the reaction mixture. Control experiments
with the solutions of manganese(III), manganese(II), and Ala
were also performed under the same experimental conditions,
which did not show any precipitation.
17]
[
16,17]
surface, is shown next
:
ꢁ
E(V vs.SCE) = E o + (2.303RT/nF) log([Mn(III)]/[Mn(II)]) [2]
−
1
−1
where R = 8.314 J K mol , T = 298 K, n = 1, and F = 96,500
−1
C mol . The attainment of redox potential of the equilibrium
in acetic acid medium required 20–30 min. Plots of E(V vs. SCE)
versus log ([Mn(III)]/[Mn(II)]) were linear with an intercept
ꢁ
ꢁ
ꢁ
of Eo – Ecal . Since Ecal = 0.246 V for the saturated calomel
ꢁ
RESULTS AND DISCUSSION
electrode (SCE) at 298 K, the formal redox potential (Eo ) of
the Mn(III)–Mn(II) couple could be calculated. The potential
measurements were made at various concentrations of acetic
acid and in the presence of added sodium salts containing ions,
Reaction Stoichiometry
The reaction stiochiometry of 2:1, observed in the prelimi-
nary studies, is represented by:
−
−
−
4−
−
−
HSO4 , and ClO4 , OAc , P2O7 , and F and Cl .
3
+
2
Mn + CH3CH(NH )COOH +H2O
2
Kinetic Measurements
(oxidant)
(Ala)
Solutions containing requisite amounts of Ala (enough to
maintain a large mole excess over Mn(III)), acetic acid (to main-
2+
+
+
→
CH3CHO +2Mn + NH + H + CO2
[3]
4
(aldehyde)

KINETICS OF ALANINE OXIDATION BY MANGANESE(III)
1333
Kinetic Orders
TABLE 1
Effects of varying concentrations of Ala, Mn(III), Mn(II), and
acetate ions on the reaction rate, with [HOAc] = 60% (v/v)
and temperature 313 K
Under the pseudo-half-order conditions of [Ala] >>
[
Mn(III)], the concentration of Mn(III) was varied while
keeping constant all other experimental conditions such as
Ala]o, [HOAc], [NaOAc]o, [Mn(II)]o, and temperature. Plots
[
3
2
5
1
0 [Mn(III)] 10 [Ala] 10 [NaOAc] 10[Mn(II)] 10 kobs
1/2
of [Mn(III)] versus time were linear for more than 75% of
1
/2 −1
(
M)
(M)
(M)
(M)
M
s )
ꢁ
the reaction (Figure 1). Pseudo-half-order rate constants (k or
kobs) remained constant with varying [Mn(III)]o (Table 1). At 0.500
constant experimental conditions, the concentration of Ala was 1.00
varied. The reaction rate increased with increase in [Ala]o (Table 2.00
2.00
2.00
2.00
2.00
2.0
3.00
3.00
3.00
3.00
3.00
3.00
3.00
3.00
3.00
3.00
3.00
1.00
2.00
3.00
4.00
5.00
3.00
3.00
3.00
3.00
3.00
1.00
1.00
1.00
1.00
1.00
1.00
1.00
1.00
1.00
1.00
1.00
1.00
1.00
1.00
1.00
1.00
1.00
2.00
3.00
4.00
5.00
2.20
2.23
2.23
2.37
2.35
2.37
2.09
2.23
4.52
6.63
8.74
1.87
2.09
2.23
3.10
4.22
2.22
2.23
2.37
2.23
2.37
1
). A plot of ln kobs versus ln [Ala]o was linear with a slope of 3.00
.94 (Figure 2), showing a fractional-order dependence of the 4.00
0
rate on [Ala]. Similarly, at constant [Mn(III)]o, [Ala]o, [HOAc], 5.00
Mn(II)]o, and temperature, the reaction rate was found to in- 1.00
2.00
1.00
2.00
3.00
4.00
5.00
2.00
2.00
2.00
2.00
2.00
2.00
2.00
2.00
2.00
2.00
[
crease with increase in [NaOAc]o (Table 1). A plot of ln kobs 1.00
versus ln [NaOAc] was linear with a slope of 0.47, showing a 1.00
−
fractional-order dependence on [OAc ]. The kinetic measure- 1.00
ments were performed using varying perchloric acid concen- 1.00
+
trations, and the effective [H ] of each reaction mixture was 1.00
evaluated from its pH value measured using a pH meter. The 1.00
1.00
1.00
1.00
1.00
1.00
1.00
1.00
1.00
+
rate of reaction decreased with increase in [H ] (Table 2). A plot
+
of ln kobs versus ln [H ] gave a negative slope of 0.46 (Figure 3),
showing an inverse fractional-order dependence on [H ]. The
+
effect of dielectric constant (D) of the medium was studied by
varying the content of HOAc (Table 2).[ The rate decreased
with the increased HOAc content (i.e., decreased D). Varying
concentrations of Mn(OAc)2 had a negligible effect on the rate
18]
1
/2
−2
FIG. 1. Plots of [Mn(III)] vs. time: [Ala] = 4.00 × 10 M; [Mn(II)] =
−
2
ꢁ
1
.00 × 10 M; [NaOAc] = 0.300 M; [HOAc] = 60% (v/v); [Mn(III)]0 = 5.00
FIG. 2. Plot of ln k vs. ln [Alanine] for the reaction, with [Mn(III)]0 = 1.00
−
3
−3
−3
−3
−3
−2
×
10 M (a), 1.00 × 10 M (b), 2.00 × 10 M (c), 4.00 × 10 M (d), and
× 10 M; [Mn(II)] = 1.00 ×10 M; [NaOAc] = 0.300 M; [HOAc] = 60%
−
3
5
.00 × 10 M (e); temperature 313 K.
(v/v); temperature 313 K (color figure available online).

1334
C. S. CHIDAN KUMAR ET AL.
TABLE 2
Effects of varying HOAc content and HClO4 concentrations on
−
3
the reaction rate, with [Mn(III)] = 1.00 × 10 M; [Ala] =
−2
−2
4
.00 × 10 M; [Mn(II)] = 1.00 × 10 M; [NaOAc] = 0.300
M; and temperature 313 K
5
1/2
5
Percent of 10 kobs(M
HOAc(v/v)
10 kobs
s ) 10 [HClO4] (M)a pH
−
1
2
(M s )
1/2 −1
50
55
60
65
70
2.50
2.31
2.23
2.10
2.00
0.0
1.0
5.0
7.5
10.0
2.68
2.66
2.64
2.59
2.56
2.27
1.60
6.9
6.0
5.9
5.6
5.4
4.3
2.0
2
3
0.0
0.0
ꢁ
−
−3
a[HClO
FIG. 3. Plot of log k vs. pH for the reaction, with [Mn(III)] = 1.00 × 10
4
] was varied at constant 60% HOAc.
2
−2
M; [Ala] = 4.00 × 10 M; [Mn(II)] = 1.00 × 10 M; [NaOAc] = 0.300 M;
[
HOAc] = 60% (v/v); temperature 313 K (color figure available online).
(
Table 1). Added chloride and perchlorate ions (in the form of
sodium salts) increased the reaction rate, while the fluoride ion
decreased it (Table 3). These effects may be attributed to the
ligand substitution of acetate ions in the complex by the added
Cl , ClO4 , or F ligand, resulting in a new Mn(III)–Mn(II)
couple with a changed redox potential, in each case. The effect
of temperature was studied in the range, 308–328 K and the ac-
tivation parameters were calculated using Arrhenius and Eyring
plots (Table 4).
served kinetic results include a fractional-order dependence of
the rate each on [Mn(III)], [Ala], and [NaOAc], and an inverse
fractional-order dependence on [H ]. An explanation for the
+
−
−
−
half-order dependence on [Mn(III)] is as follows: The occur-
rence of a fast preequilibrium showing the formation of the
−
kinetically active complex species [Mn(OAc)4] is supported
by the strong tendency of Mn(III) acetate in the presence of
excess of acetate ions to form coordination compounds. The
equilibrium accounts for the fractional-order dependencies on
Spectroscopic Evidence for the Intermediate Complex
The visible spectrum of manganese(III) acetate solution both [Mn(III)] and [acetate ion]. Amino acids (S) in strongly
+
shows a λmax at 470 nm. When Ala reacts with manganese(III), acidic medium exist in their protonated SH form, whose dis-
the maximum shifts to 446 nm, suggesting the formation of an sociation in the first fast equilibrium accounts for the negative
+
Ala–Mn(III) complex species.
fractional-order dependence on [H ].
Scheme 1 shows a detailed electronic interpretation of the
Solvent Isotope Studies
Ala oxidation mechanism. The reactive oxidant species, a
−
The reaction was studied in D2O solvent (0.4 at% D) to de- four-coordiante complex X [Mn(OAc)4] , is formed in a fast
[
19]
termine the isotope effect. The rate constants under identical preequilibrium (step ii). The Mn(III) ion is known to form four-
−5
1/2 −1
[12,13]
conditions were kH = 1.95 × 10
M
s
and kD = 8.90 coordinate square planar complexes similar to X.
Com-
−
6
1/2 −1
×
10 M s , resulting in the solvent primary isotope effect plex X interacts with the substrate (Ala) in the rate-determining
ꢁ
(
kH/kD) of 2.2.
step (iii) to form the intermediate Mn(III)–Ala complex X ,
In other studies, it was proposed on the basis of spectropho- where Ala as a bidentate ligand coordinates through its amino
tometric and electrochemical experiments that the reactive man- N atom and carboxylic O atom. Fast step (iv) involves single-
ganese(III) species in aqueous acetic acid-acetate medium was electron transfers (SETs) and other rearrangements to form
−
[18,19]
ꢁꢁ
[
Mn(OAc)3] or [Mn(OAc)4] .
In the present study, the ob- an amino alkyl carbon free radical X , consistent with the
TABLE 3
−
−
−2
−
−3
Effects of varying concentrations of Cl , F , and ClO4 ions on the reaction rate, with [Mn(III)] = 1.00 × 10 M; [Ala] = 4.00
−2
×
10 M; [Mn(II)] = 1.00 × 10 M; [NaOAc] = 0.300 M; [HOAc] = 60% (v/v); and temperature 313 K
2
−
5
1/2 −1
2
−
5
1/2 −1
2
−
5
1/2 −1
1
0 [Cl ] (M)
10 kobs (M s )
10 [F ] (M)
10 kobs (M s )
10 [ClO4 ] (M)
10 kobs (M s )
0.5
1.0
1.5
2.10
2.23
2.89
0.5
1.0
1.5
2.50
2.23
2.10
0.5
1.0
1.5
2.09
2.23
2.74

KINETICS OF ALANINE OXIDATION BY MANGANESE(III)
1335
SCH. 1. Mechanism of alanine oxidation by Mn(III) in acid solutions.
positive free radical test, along with MnII acetate and car- nia to produce the end products, acetaldehyde and ammonium
bon dioxide. In fast step (v), the second Mn(III) ion picks ion.[
21,22]
ꢁ
ꢁ
up an electon from the carbon radical, X , to form Mn(II)
The rate of Ala oxidation is given by the slow step in Scheme
and amino carbocation. In step (vi), a nucleophilic attack by 1 as
H2O molecule on the electrophilic carbon center leads to the
formation of a protonated amino ethanol, which then forms
−
Rate = −d[Mn(III)/dt = k3 [Mn(OAc) ][S]
[4]
4
a protonated aldehyde and ammonia in step (vii). In the final After the application of the steady-state concept to the interme-
−
step, the protonated acetaldehyde is deprotonated by ammo- diate complex, [Mn(OAc)4 ], and the substitution for [S] from

1336
C. S. CHIDAN KUMAR ET AL.
equilibrium (i), Eq. (5) is obtained:
added electrolyte, respectively. According to Eq. (8), a plot of
ꢁ
1/2
log k versus µ should be linear with a slope of 1.02ZAZB and
intercept of log ko. As the slope of the line depends on ZAZB,
the reaction rate may increase, decrease, or remain unaffected
with the ionic strength of the medium. In the present case, the
zero effect of ionic strength on the rate signifies that at least one
of the reactants is a neutral molecule as shown in step (iii) of
−
−
+
+
[
Mn(OAc) ] = k2 [Mn(OAc) ][OAc ][H ]/{k [H ]
4
3
−2
+
+
K1k3[SH ]}
[5]
A similar treatment of S along with the use of equilibrium step
−
(
ii) for [Mn(OAc)4 ] results in:
Scheme 1. Hence, the observed ionic strength effect is consis-
tent with the Bronsted–Bjerrum concept[
scheme.
23,24]
for the proposed
+
+
+
+
+
[
S] = k1[SH ]{k [H ] + K1k3[SH ]}/{k [H ]{k [H ]
−
2
+
−1
−2
+
−
+
K1k3[SH ]} + k2k3[Mn(OAc) ][H ][OAc ]}
[6]
3
A negative dielectric constant (D) effect, determined by vary-
ing [HOAc], also provides support for the proposed mecha-
nism (Scheme 1). The effects of solvents on the reaction ki-
−
Substitutions for [Mn(OAc)4 ] and [S] from Eqs. (5) and
(6), respectively, in Eq. (4) lead to the rate law [Eq. (7)]
[25]
[26]
netics have been described by House, Entelis and Tiger,
shown next.
[23]
[27]
Amis, and Richardt. For a limiting case of zero angle of
approach between two dipoles or an ion–dipole system, Amis
has shown that a plot of log k versus 1/D is linear with a neg-
[
23]
+
−
+
Rate = K1k2k3[SH ][Mn(OAc) ][OAc ]/[K k3[SH ]
ꢁ
3
1
+
−
+
[H ]{k + (k k3/k )[Mn(OAc) ][OAc ]} [7] ative slope for interactions between a negative ion and a dipole
−2
2
−1
3
or between two diploes, while a positive slope results for the
The observed kinetic data of fractional-order dependencies of positive ion–dipole interactions. The negative dielectric effect
+
the rate on [Ala] (or [SH ] in acid medium), [Mn(III)] and observed in the present study clearly supports the anion–dipole
[
−
+
OAc ], and a negative fractional-order on [H ] are supportive interactions in the slow step of Scheme 1. Solvent isotope studies
of the rate law (Eq. (7)), which has been derived based on the have shown that the rate of reaction is higher in D2O medium for
proposed mechanistic Scheme 1.
The zero-effect of the ionic strength and the negative effect fast preequilibrium with H3O or OH ion transfer, the rate in-
of dielectric constant (D) indicate the involvement of neutral creases in D2O medium since D3O and OD are a stronger acid
and negative species in the rate-determining step in Scheme 1. and stronger base, respectively, than H3O and OH ions. In the
The reduction product, Mn(II), does not influence the rate as it present case, where the reaction has negative fractional order in
is not involved in a preequilibrium.
[
23,28]
acid or base catalyzed reactions.
For a reaction involving a
+
−
+
−
+
−
+
ꢁ
ꢁ
[H ], the observed solvent isotope effect of k (H2O)/k (D2O) >
+
The proposed reaction mechanism and the derived rate law 1 is due to the greater acid effect of the D ion as compared to
+
are also supported by the following experimental facts:
the H ion. The magnitude of the effect in D2O is consistent
[
23]
Bronsted and Bjerrum
have explained the primary salt with the expected value of two to three times greater, which is
effect on the reaction rates through the relation
supporting the negative fractional-order dependence of the rate
on [H ]. The moderate values of energy of activation and other
[8] activation parameters are supportive of the proposed mecha-
nism. The fairly positive values of free energy of activation and
+
ꢁ
1/2
log k = log ko + 1.02 ZAZB µ
where µ is the ionic strength of the medium, A and B are the enthalpy of activation indicate that the transition state is highly
reacting ions, ZA and ZB are charges on the respective species, solvated, while the negative entropy of activation suggests an as-
ꢁ
k and ko are rate constants in the presence and absence of the sociative, rigid transition state with reduced degrees of freedom.
ꢂ
Furthermore, the negative ꢂS suggests that the Ala–Mn(III)
redox reaction is controlled by entropy rather than enthalpy.
TABLE 4
Kinetic and activation parameters for the reaction, with
−
3
−2
[
=
Mn(III)] = 1.00 × 10 M; [Ala] = 4.00 × 10 M; [Mn(II)]
−2
CONCLUSIONS
1.00 × 10 M; [NaOAc] = 0.300 M; [HOAc] = 60% (v/v);
The Mn(III)–Ala reaction has been investigated in acid so-
lutions. The reaction stoichiometry involving the Ala oxidation
to acetaldehyde has been found to be 2:1.
and temperature 313 K
Temperature (K) 10 kobs (M1/2 s )
5
−1
The redox reaction has the following experimental rate
3
3
3
3
3
08
13
18
23
28
1.82
2.23
3.63
6.96
10.72
log A
Ea
ꢂH
1.65
59.8 kJ mol
ꢁ
0.5
0.9
− 0.5
+ 0.5
law: Rate = k [Mn(III)] [Ala] [OAc ] /[H ] , where each
species shows a fractional order. A suitable mechanism consis-
tent with the experimental observations has been proposed to
derive a rate law.
−
−1
1
ꢂ
69.7 kJ mol
ꢂ
−1
−1
ꢂS −192 JK mol
ꢂ
−1
ꢂG
61.6 kJ mol
ꢂ
ꢂ
ꢂ=
The activation parameters, Ea, ꢂH , ꢂG , and ꢂS , have
aMean value.
been determined to understand whether the reaction is controlled

KINETICS OF ALANINE OXIDATION BY MANGANESE(III)
1337
ꢂ
−1
−1
by entropy or enthalpy. The negative ꢂS (−192 J K mol ) 15. Chandraju, S.; Rangappa, K.S.; Mahadevappa, D.S.; Made Gowda, N.M.
Synth. React. Inorg. Metal-Org. Chem. 1997, 27(9), 1329.
indicates the rigid transition state formation and the entropy-
1
6. Vogel, A.I. Qualitative Organic Analysis; Longman and Green: London,
958, pp. 708–720.
7. Chandraju, S.; Rangappa, K.S.; Made Gowda, N.M. Int. J. Chem. Kinet.
controlled reaction.
1
1
1
2
998, 30007. (b) Meenakshisundaram, S.; Rajagopal, V. Coord. Chem. Acta
003, 76(1), 75. (c) Goswamy, K.B.; Chandra, G.; Srivastava, S.N. J. Indian
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3

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