Kinetics and mechanisms for TCE oxidation by permanganate

Article abstract of DOI:10.1021/es991279q

The oxidation of trichloroethylene (TCE) by permanganate proceeds in three sequential reaction steps. In the initial step, a cyclic hypomanganate ester is formed via an activated organometallic complex. The activation parameters for this step were determined to be E(a) = 41.46 k J/mol, ΔH(+) = 39 k J/mol, and ΔS(+) = -14 J/mol. The initial reaction is a rate-limiting step (second-order rate constant k(1p) = 0.65-0.68 M^-1 s^-1 at 21 °C) and independent of pH. In the second step, the decomposition of the cyclic ester with complete chlorine liberation proceeds quickly via various reaction pathways to form four carboxylic acids. Approximately 77% of the TCE was transformed to formic acid at pH 4, while 95-97% of the TCE was transformed to oxalic and glyoxylic acids at pH values of 6-8. Kinetic data suggest that the decomposition rate of the cyclic ester is at least 100 times higher than its formation rate. In the final step, all carboxylic acids are oxidized by permanganate to the final product, CO₂. Second-order rate constants of k(3ap) = 0.075-0.35 M^-1 s^-1, k(3bp) = 0.13-0.37 M^-1 s^-1, and k(3cp) = 0.073-0.11 M^-1 s^-1 over a pH range of 4-8 at 21 °C were estimated for oxidation of formic, glyoxylic, and oxalic acids, respectively. The oxidation rate of carboxylic acids and accumulation rate of CO₂ increase with decreasing pH. The kinetic model that was developed, formulated, and solved analytically on the basis of the understanding of various processes is consistent with results obtained in the kinetic experiments. The oxidation of trichloroethylene (TCE) by permanganate proceeds in three sequential reaction steps. In the initial step, a cyclic hypomanganate ester is formed via an activated organometallic complex. The activation parameters for this step were determined to be E_a = 41.46 kJ/mol, ΔH^? = 39 kJ/mol, and ΔS^? = -14 J/mol. The initial reaction is a rate-limiting step (second-order rate constant k_1p = 0.65-0.68 M^-1 s^-1 at 21°C) and independent of pH. In the second step, the decomposition of the cyclic ester with complete chlorine liberation proceeds quickly via various reaction pathways to form four carboxylic acids. Approximately 77% of the TCE was transformed to formic acid at pH 4, while 95-97% of the TCE was transformed to oxalic and glyoxylic acids at pH values of 6-8. Kinetic data suggest that the decomposition rate of the cyclic ester is at least 100 times higher than its formation rate. In the final step, all carboxylic acids are oxidized by permanganate to the final product, CO₂. Second-order rate constants of k_3ap = 0.075-0.35 M^-1 s^-1, k_3bp = 0.13-0.37 M^-1 s^-1, and k_3cp = 0.073-0.11 M^-1 s^-1 over a pH range of 4-8 at 21°C were estimated for oxidation of formic, glyoxylic, and oxalic acids, respectively. The oxidation rate of carboxylic acids and accumulation rate of CO₂ increase with decreasing pH. The kinetic model that was developed, formulated, and solved analytically on the basis of the understanding of various processes is consistent with results obtained in the kinetic experiments.

Full text of DOI:10.1021/es991279q

Environ. Sci. Technol. 2000, 34, 2535-2541
1 2 3
γ and η are stoichiometric coefficients, and k , k , and k are
Kinetics and Mechanisms for TCE
Oxidation by Permanganate
rate constants. In the first reaction, TCE disappearance was
found to be independent of pH. Complete chlorine liberation
was thought to occur in the second step of eq 1. However,
the second and third reactions have not been studied in
detail. This study was undertaken to understand fully the
kinetics and mechanisms for TCE oxidation by permanganate
and more specifically (i) to identify reaction products, (ii) to
elucidate reaction pathways and their pH dependence, (iii)
to develop a kinetic model, and (iv) to determine the rate
Y. E U G E N E YA N * , † A N D
F R A N K L I N W . S C H W A R T Z
Department of Geological Sciences, The Ohio State University,
Columbus, Ohio 43210
constants for all major reactions in the KMnO
system.
4
2
-TCE-H O
The oxidation of trichloroethylene (TCE) by permanganate
proceeds in three sequential reaction steps. In the
initial step, a cyclic hypomanganate ester is formed via
an activated organometallic complex. The activation
parameters for this step were determined to be Ea ) 41.46
Experimental Methods
1
4
Materials. Radiolabeled [1,2- C]TCE (3.1 mCi/ mmol) was
purchased from Sigma Chemical Co. (St. Louis, MO) and
diluted with methanol in a 10-mL glass volumetric flask and
stored at 4 °C in the dark. Carboxylic acids were obtained
from Fluka (Buchs, Switzerland) with purity 99+% and used
as received. All other chemicals were of the highest purity
available and were described in the previous study (7).
Product Studies and Kinetic Experim ents. With TCE
oxidation by permanganate, carboxylic acids and carbon
dioxide would be expected as major intermediate and final
products (7). Two types of kinetic tests were designed to
q
q
kJ/mol, ∆H ) 39 kJ/mol, and ∆S ) -14 J/mol. The
initial reaction is a rate-limiting step (second-order rate
-
1 -1
constant k1p ) 0.65-0.68 M
s
at 21 °C) and independent
of pH. In the second step, the decomposition of the
cyclic ester with complete chlorine liberation proceeds
quickly via various reaction pathways to form four carboxylic
acids. Approximately 77% of the TCE was transformed to
formic acid at pH 4, while 95-97% of the TCE was transformed
to oxalic and glyoxylic acids at pH values of 6-8. Kinetic
data suggest that the decomposition rate of the cyclic
ester is at least 100 times higher than its formation rate.
In the final step, all carboxylic acids are oxidized by
permanganate to the final product, CO2. Second-order
2
identify the carboxylic acids and to trace CO formation. Each
type consists of three experiments conducted at pH values
of 4, 6, and 8 at 21 °C for about 8 h, and each experiment
has one set of vials with two replicates and one control vial
containing only TCE aqueous solution.
For investigating carboxylic acid distributions over time,
a first kinetic test was conducted in 50-mL glass vials with
PTFE/ silicone septum-lined screw-top caps. Each vial was
filled with a 50-ml TCE solution of 0.1 mM. The experiment
-
1
-1
rate constants of k3ap ) 0.075-0.35 M
s
, k3bp ) 0.13-
-
1
-1
, and k3cp ) 0.073-0.11 M s-1 over a pH
-1
0
.37 M
s
range of 4-8at 21 °C were estimated for oxidation of formic,
glyoxylic, and oxalic acids, respectively. The oxidation
rate of carboxylic acids and accumulation rate of CO2
increase with decreasing pH. The kinetic model that was
developed, formulated, and solved analytically on the
basis of the understanding of various processes is consistent
with results obtained in the kinetic experiments.
4
began by injecting 5 mL of 6.3 mM KMnO through the
septum. A second needle enabled an equal volume of TCE
solution to be displaced. Vials were placed on a vibratory
shaker until sampled. During sampling for each experiment,
two vials were taken. A 1-mL aliquot of reaction solution
from each vial was transferred to a volumetric flask with
appropriate dilution following immediate TCE analysis. To
quench the reaction, 1 mL of thiosulfate from the stock
solution was added to each vial. The quenched solution was
centrifuged at 3000 rpm for 20 min and filtered by a 0.2-µm
filter to separate the precipitated manganese dioxide. After
filtration, the solutions were withdrawn for carboxylic acid
analysis.
Introduction
There is considerable interest in the in-situ oxidative
degradation of chlorinated ethenes with a metal-oxo reagent
(
1-5). In terms of remediating contaminated groundwater,
this method is attractive due to the rapid degradation of
chlorinated ethenes, the nonreactive nature of the reagent
with carbonate and bicarbonate, the ease of field imple-
mentation, and the relatively low costs (1, 5, 6). Degradation
kinetics for five chlorinated ethenes, including PCE, TCE,
and three isomers of DCE were reported by ref 7. TCE
degradation was postulated to involve the following sequen-
tial reactions:
To trace the CO production, a second kinetic test was
2
undertaken with radiolabeled 14C TCE. To maintain the same
initial conditions as the first kinetic test, experiments began
with the addition of 0.1 mL of permanganate stock solution
to a 3-mL test tube filled with 2.9 mL of 0.093 mM TCE
solution. Each experiment consists of one set of test tubes
with duplicates. At predetermined interval, the reaction was
quenched by adding 0.1 mL of thiosulfate from the stock
k1
98 I
k2
-
-
2
solution. The test solution was then analyzed for CO .
C HCl + MnO
98 γCA + MnO + 3Cl
(1)
(2)
2
3
4
2
In addition, five sets of duplicated kinetic experiments
were performed at various temperatures from 5 to 25 °C to
investigate how temperature affects the rate of TCE loss and
to determine activation parameters for the initial oxidation.
This temperature range covers variation commonly observed
for natural groundwaters. All experiments were conducted
by reacting 0.06 mM TCE with an excess of MnO
at pH 7 and ionic strength of 0.05 M.
k3(MnO -)
4
γCA
8 ηCO2
where I is a cyclic complex, CA are various carboxylic acids,
-
4
(1 mM)
*
Corresponding author telephone: (630)252-6322; fax: (630)-
2
52-5747; e-mail: eyan@anl.gov.
†
Chem ical Analysis. The organic acids were analyzed on
a Waters high-performance liquid chromatography (HPLC)
Present address: Argonne National Laboratory, 9700 S. Cass Ave.,
Argonne, IL 60439.
1
0.1021/es991279q CCC: $19.00

2000 Am erican Chem ical Society
VOL. 34, NO. 12, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
2 5 3 5
Published on Web 05/18/2000

fitted with a Bio-Rad Aminex HPX-87H column (300 × 7.8
mm i.d.). The column was operated at room temperature
using degassed dilute sulfuric acid (4 mM) as eluent at a flow
rate of 0.6 mL/ min. The column effluents were monitored
by an UV spectrophotometric detector (Waters 486) at 210
nm. The products were identified and quantified by com-
Kinetic Equations for TCE Oxidation. TCE oxidation by
permanganate ion generally involves the formation and
decomposition of cyclic hypomanganate ester (CHME) and
the oxidation of carboxylic acids, based on the TCE trans-
formation pathways in Figure 1. A second-order reaction for
the formation of CHME was suggested by (7)
parison with external standards. CO
2
was determined by
under
k1p
retaining CO under basic conditions and stripping CO
2
2
-
C HCl + MnO
4
8CHME
(3)
2
3
acidic conditions. This method was based on and adapted
from ref 8. TCE concentrations were measured as described
in ref 7.
where k1p is a second-order rate constant. The decomposition
of CHME via various pathways to form major carboxylic acids
and the release of chlorides can be described by eqs 4-6:
Results and Discussion
k2a
Products. In the experiments for identifying products, the
reaction was initiated with 0.09 mM TCE reacting with 0.63
mM permanganate in a phosphate-buffered solution having
an ionic strength of 0.05 M at pH 4, 6, and 8, respectively.
Four carboxylic acids (formic, oxalic, glyoxylic, and glycolic
acids) were identified in the system as intermediate products.
Either formic acid or oxalic acid predominated, depending
on pH. At 1 h after the reaction started, a maximum of 43%
of the initial TCE was converted to either formic acid or oxalic
acid, and up to 25% of the TCE was transformed to glyoxylic
acid. Glycolic acid was measured in very small quantities
RCHME + âH O 8 2RHCOOH +
2
-
+
RMnO + 3RCl + γH (4)
2
k2b
RCHME + ηH O 8 ROHC-COOH +
2
-
+
RMnO + 3RCl + νH (5)
2
k2c
RCHME + êH O 8 RHOOC-COOH +
2
-
+
RMnO + 3RCl + úH (6)
2
(
<2%) in the reaction.
The radiolabled product analysis indicates that the
where R, â, γ, η, ν, ê, and ú are stoichiometric coefficients
and k2a, k2b, and k2c are rate constants for the formation of
formic, glyoxylic, and oxalic acids, respectively.
majority of the TCE was eventually transformed to CO
final product. Assuming that TCE is stoichiometrically
converted to CO , 57-88% of the initial TCE at various pH
2
, the
2
The oxidation of formic, glyoxylic, and oxalic acids by
permanganate involves a series of reactions. Most studies,
however, have shown that the reaction at a rate-limiting step
can be described by second-order reaction for formic and
oxalic acids (17-19). By analogy with the oxidation of formic
and oxalic acids, it is reasonable to assume a second-order
reaction for glyoxylic acid because it has a structure similar
to the other two acids. The rate laws for these three oxidation
reactions are shown in eqs 7-9 (Table 1). With an excess of
values from 4 to 8 was converted to CO
were terminated at 8 h.
2
when the experiments
The proposed chemical pathways for TCE oxidation
Figure 1) were developed from a basic understanding of
(
alkene oxidation by permanganate in organic synthesis
chemistry (9-13). The products and intermediates in the
shadowed boxes were identified in our experiments. TCE
oxidation is initiated by the attack of permanganate ion, as
an electrophile, on the carbon-carbon double bond (7). An
intermediate, cyclic hypomanganate ester 2, is formed via
an activated organometallic complex (12, 14, 15). The rapid
decomposition of cyclic ester 2 can follow several different
pathways in an aqueous system such as oxidative hydrolysis
and hydrolysis.
-
4
in aqueous media at a fixed pH, the order of kinetic
MnO
reaction can be reduced. The rates of disappearance/
appearance for TCE, major intermediates, and final product
are summarized in Table 1, eqs 10-15.
Estim ation of Rate Constants. To find the rate constants
for the transformation and/ or formation of TCE, three
Oxidative hydrolysis transforms cyclic hypomanganate
ester 2 rapidly through disproportionation (13) to cyclic
manganate(VI) ester 3 followed by subsequent fragmentation
of 3 to form formic acid 5 via 4. Two other possible pathways
involve the hydrolysis of cyclic ester 2 to acyclic hypoman-
ganate(V) ester 6. The acyclic manganate(VI) ester 7, oxi-
datively hydrolyzed from 6, may undergo rapid electron
transfer and then either hydrolysis to form glycolic acid 9 via
2
carboxylic acids, and CO , eqs 10-15 (Table 1) were solved
analytically using Laplace transformation. The solutions for
TCE, formic, glyoxylic, and oxalic acids in terms of normalized
concentration are
[TCE]
k1t
)
e-
(16)
[TCE]0
8
or release of one hydrogen chloride and hydrolysis to form
e-kit
[
C ]
t l
i, j ) 1, 2, and 3l;
[l ) a, b, and c;
glyoxylic acid 12 or oxalic acid 13 via 11. In weak alkalinic
to alkalinic solutions, acyclic ester 6 may simply hydrolyze
to trichloroglycol 10, which rapidly releases two hydrogen
chlorides to form 11. The compound 11 is subsequently
hydrolyzed to either glyoxylic acid 12 or oxalic acid 13. The
intermediates, 4, 8, and 11, are all acyl halides that are so
reactive that hydrolysis is easily carried out to form carboxylic
acids (16). Carboxylic acids are further oxidized to carbon
dioxide at a relatively slow rate.
)
mk1k2l∑i
]
[TCE]0
(
k - k )
j i
∏
j(j*i)
2
l ) a
m )
; (17)
[
]
1 l ) b, c
where [C
t
]
l
is the concentration of the carboxylic acids and
the subscript l ) a, b, and c denotes formic, glyoxylic, and
oxalic acids, respectively. The formation of CO
described as
2
can be
The liberation of chlorine substitutes is hypothesized as
a series of reactions, involving the transformation of hy-
pomanganate ester to manganate ester, the electron transfer
to release hydrogen chloride, and hydrolysis. A stoichio-
metrically consistent release of chloride was observed along
with disappearance of TCE (7). Thus, the decomposition of
the ester with chloride release is a very rapid process, which
is much faster than its formation.
c
e-kit
[
CO2]
)
2 1 - k1
k k
i, j )
∑
2l 3l∑i
[
TCE]0
l)a
ki
(k - k )
j
i
(
∏
)
{
}
j(j*i)
1, 2, and 3l (18)
2
5 3 6
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 34, NO. 12, 2000

FIGURE 1. Proposed TCE oxidation pathways. Shadowed boxes are products identified in this study.
To establish a kinetic model, seven rate constants (k
2b, k2c, k3a, k3b, and k3c) need to be estimated using kinetic
data from the TCE oxidation experiments. Accordingly, the
kinetic experiments described previously were conducted
1
, k2a
,
cyclic hypomanganate ester decomposition along various
pathways. Thus, our kinetic model for TCE oxidation is
consistent with experimental results at pH ranging from 4
to 8.
k
with a H
3
PO
4
-K
2
HPO
4
-KH
2
PO
4
-NaOH buffer at pH values
The estimated rate constants are listed in Table 2. With
of 4, 6, and 8 at 21 °C. The concentrations of TCE, the
intermediates (including formic, glyoxylic, and oxalic acids),
and the final product (CO ) were measured with time. Mass
2
the pseudo-first-order rate constant k
1
ranging from 4.11 ×
, the second-order rate constant k1p for
the initial TCE oxidation, defined in eq 10, can be calculated.
-
4
-4 -1
10 to 4.30 × 10
s
-
1
-1
balances in these experiments ranged from 0.84 to 1.09. The
results indicate that the experiments were well-controlled,
even though various analytical procedures were involved.
To estimate rate constants, eqs 16 and 17 were used to
fit the measurements, which are plotted as discrete data
points in Figure 2. The fitted curves are shown as solid lines.
The results show good agreement across the experiments at
three different pH levels in terms of TCE destruction and
The calculated k1p in the range of 0.65-0.68 M
s
agrees
-
1
-1
with ref 7 (i.e., k ) 0.67 ( 0.03 M
s
over a pH range of
-
1
4-8). A similar value of rate constant (k ) 0.89 ( 0.03 M
-
1
s
at pH 7 and 20 °C) was recently reported in ref 20.
To choose an appropriate k for the second step in TCE
oxidation, the concentration data for the carboxylic acids
were fitted with eq 17 using various ratios of k to k . The
results indicate that determination coefficient (r ) for curve
2
2
1
2
VOL. 34, NO. 12, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
2 5 3 7

TABLE 1. Listing of Kinetic Equationsa
equation
no.
d[HCOOH]t
-k3ap[HCOOH] [MnO4-]
)
t
7
dt
d[OHC-COOH]t
-k3bp[OHC - COOH] [MnO4-]
)
t
8
9
dt
d[HOOC-COOH]t
-k3cp[HOOC - COOH] [MnO4-]
)
t
dt
d[C HCl ]
2
3
-k [C HCl ] k1 ) k1p[MnO -]
)
)
1
2
3
4
10
11
12
13
14
15
dt
d[CHME]
dt
k [C HCl ] - k [CHME] k2 ) k + k2b + k2c
1
2
3
2
2a
d[HCOOH]t
dt
k3a ) k3ap[MnO4-]
)
2k2a[CHME] - k3a[HCOOH]t
d[OHC - COOH)t
k3b ) k3bp[MnO4-]
)
k2b[CHME] - k3b[OHC - COOH]t
dt
d[HOOC - COOH]t
k3c ) k3cp[MnO4-]
)
k [CHME] - k [HOOC - COOH]
2c 3c
t
dt
d[CO2]
)
k3a[COOH] + 2k [OHC - COOH] + 2k [HOOC-COOH]
t 3b t 3c t
dt
+
d[HCOOH]t
dt
k3ap1[H ] + k3ap2Ka
-
-
[HCOOH] [MnO4-]
t
)
-(k3ap1[HCOOH] + k3ap2[HCOO ])[MnO4 ] ) -
(
+
)
25
26
[
H ] + K
a
d[OHC - COOH]t
-
-
)
-(k3bp1[OHC - COOH] + k3bp2[OHC - COO ])[MnO ] )
4
dt
+
k3bp1[H ] + k3bp2Ka
-
[OHC - COOH] [MnO4-]
(
+
)
t
[
H ] + Kb
d[HOOC - COOH]t
-
-
-
-
)
-(k3cp1[HOOC - COOH] + k
[HOOC - COO ] + k [ OOC - COO ])[MnO ] )
3cp2
3cp3
4
dt
+
2
+
k3cp1[H ) + k
K [H ] + k K K
27
3cp2 c1
cp3 c1 c2
[HOOC - COOH] [MnO4-]
-
t
(
+
+ 2
)
K K + K [H ] + [H ]
c1 c2
c1
a
Subscript t for each acid denotes the total concentration; k
1
, k3a, k3b, and k3c are pseudo-first-order rate constants; k2a, k2b, and k2c are first-order
rate constants; k3ap, k3bp, and k3cp are second-order rate constants; k3ap1, k3ap2, k3bp1, k3bp2, k3cp1, k3cp2, and k3cp3 are rate constants.
fittings was improved with increasing k
.9 at k / k > 10. The large k / k ratio supports the idea that
the decomposition of cyclic ester is much faster than its
2
/ k
1
ratio and reaches
in the final step. Compared to the other steps in TCE
oxidation, the oxidation rates for the carboxylic acids are
slower than any of the rates in the previous steps. The second-
order rate constants, k3ap, k3bp, and k3cp, can be calculated
with eqs 12-14 (Table 1). In this final step, the oxidation rate
0
2
1
2
1
2
formation, as suggested by refs 7 and 9. When k
2
g 10 k
1
,
2
r
does not show any significant improvement (the increase
2
2
5
of r is less than 0.001 from k
rate constant ratios (k2a/ k , k2b/ k
step and rate constants (k3a, k3b, and k3c) for the third step
remain constant with increasing k / k . Thus, the TCE
2
) 10 k
1
to k
2
) 10 k
1
). The
of carboxylic acids (k
3
) might be underestimated by eqs 12-
-
2
2
, and k2c/ k
2
) for the second
14 because of relatively low initial concentration of MnO
4
used in the experiments. However, the estimated values for
3ap, k3bp, and k3cp are comparable with results from other
studies. For example, k3ap for the oxidation of formic acid
2
1
k
oxidation process involving TCE decomposition and car-
boxylic acid formation and decomposition is significantly
-
1
-1
ranging from 0.075 to 0.35 M
s
is consistent with the
-1 -1
controlled by the initial step (k
regardless of the rate (k ) for the second step when k
was assumed in all curve fittings.
, k2b/ k , and k2c/ k represent the fraction
1
), a rate-limiting step,
range of 0.003-0.25 M
3cp in the range of 0.073-0.11 M
for oxalic acid is also reasonable if it is compared to k )
s
obtained in ref 21. The estimated
g 102
k
-1
s
-1
at pH values of 4-8
2
2
2
k
1
. Thus, k
2
) 10 k
1
-
1
-1
The ratios k2a/ k
2
2
2
0.47-0.58 M
s , which is obtained at relatively low pH
(pH 1.9) without the initial presence of manganese(II) species
of the total decomposition rate of the cyclic ester along each
pathway and quantify the contributions from the various
pathways in Figure 1. Their values (Table 2) suggest that the
pathway leading to formic acid or oxalic acid is dominant in
the decomposition process.
(19). The rate constant k3bp for glyoxylic acid was estimated
-
1
-1
to be 0.13-0.37 M
s . On the basis of estimated rate
constants, the oxidation rate of formic and glyoxylic acids is
faster than oxalic acid over a pH range of 4-8.
The pseudo-first-order rate constants, k3a, k3b, and k3c
,
Using all the estimated rate constants, the calculated
describe the oxidation of formic, glyoxylic, and oxalic acids
2
accumulation of CO is plotted in Figure 2. The results for
2
5 3 8
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 34, NO. 12, 2000

CO
except at pH 8 in which the predicted CO
observed at later time. This overestimation may result from
the adsorption of carboxylic acids on MnO , which is more
2
are in general agreement with the experimental data,
2
is higher than the
2
readily precipitated at pH 8 than at other lower pH conditions.
Because the precipitates were left in the filtrates in our
experiments for acid analysis, acid adsorption on MnO
cause a mass deficiency, which occurs at late time (Figure
c). Consequently, the rate constant k at pH 8 may be slightly
overestimated.
Effects of pH and Tem perature. During the first step in
TCE oxidation, the rate constant k does not show strong
correlation with pH. Our previous study (7) indicated that
i) the variation of k at different pH values is small and within
the experimental error and (ii) the second-order rate constant
2
could
2
3
1
(
1
-
1
-1
k
1p is equal to 0.67 ( 0.03 M
s
over the pH range of 4-8
at 21 °C. Thus, the rate of TCE disappearance is independent
of pH. However, the fate of cyclic hypomanganate ester and
the product distribution are strongly dependent on pH.
Kinetic data demonstrate two major decomposition pathways
for cyclic hypomanganate ester. At pH 4, the transformation
of cyclic ester to formic acid is overwhelmingly dominant
(k2a/ k ) 0.77). With this relatively acidic condition, the
2
majority of the cyclic ester has been cleaved at the carbon-
carbon double bond to form two formic acids 5 via 4 (Figure
1
). At pH 6-8, decomposition favors the formation of oxalic
and glyoxylic acids ((k2b + k2c)/ k ) 0.95-0.97). A different
2
reaction pathway has been suggested for this higher pH range;
most of cyclic ester transforms to acyclic ester 6 and 7 followed
by hydrolysis to form oxalic and glyoxylic acids, 13 and 12,
via 11.
The rate constants (k3a, k3b, and k3c) for oxidation of
carboxylic acids are strongly correlated with pH over a range
of 4-8 (Figure 3). With decreasing pH, the oxidation rates
(
k
k
3
3
) increase for all three carboxylic acids. This pH effect on
results in the relatively rapid formation of CO at low pH
values as compared to higher pH values.
The temperature effect on TCE disappearance rate by
MnO was investigated by five sets of duplicated kinetic
2
4
experiments at temperatures varying from 5 to 25 °C. The
estimated logarithm of the second-order rate constant k1p
from the experiments is plotted against the reciprocal of
temperature in Figure 4. The dependence of k1p on tem-
perature, as expected, follows the Arrhenius equation. The
a
activation energy (E ) and preexponential (A) were calculated
and are listed in Figure 4. These values can provide a rate
constant, k1p, at any temperature relevant to groundwater.
7
-1
On the basis of E
a
) 41.46 kJ/ mol and A ) 1.678 × 10 M
s- , activation parameters ∆H and ∆S can be calculated.
The low activation enthalpy (∆H ) 39 kJ/ mol) obtained here
1
q
q
q
is comparable to the results from many other studies in alkene
oxidation by permanganate. This result indicates that initial
step of TCE oxidation may proceed via a similar transition
q
state as other alkene oxidation. The negative entropy (∆S )
-
14 J/ mol) is consistent with a bimolecular reaction in which
the transient complex may be highly structured and possibly
solvated.
Proposed Mechanism for TCE Oxidation
The fact that the loss of TCE is independent of pH but that
the nature and distribution of products are highly dependent
on pH strongly supports the existence of a short-lived
intermediate in the initial stage of oxidation. There has been
more than a century of study on alkene oxidation by
permanganate. Cyclic hypomanganate ester is generally
thought to form as an intermediate, on the basis of cis-
addition products, cis-diol/ ketol (22, 13), and oxygen transfer
FIGURE 2. Product distribution with time in TCE oxidation at pH
(a), pH 6 (b), and pH 8 (c). Lines for TCE and carboxylic acids rep-
resent best fits using eqs 16 and17, whereas the line for CO is cal-
culated by eq 18 using rate constants from the previous curve fitting.
-
18
4
from MnO
4
to diol observed in O-labeling experiments
2
(9). In recent studies, however, a more comprehensive
reaction mechanism has been postulated with the introduc-
VOL. 34, NO. 12, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
2 5 3 9

TABLE 2. Pseudo-First-Order Rate Constants Obtained for TCE Transformation Pathways Using Eqs 16 and 17a
k1 (10- s-1
4
)
k2a/k2
k3a (10-4 s-1
)
r 2
k2b/k2
k3b (10-4 s-1
)
r 2
k2c/k2
k3c (10-4 s-1
)
r 2
pH
4
6
8
4.30
4.11
4.11
0.77
0.02
0.04
2.19
1.53
0.47
0.98
0.91
0.97
0.03
0.32
0.42
2.34
1.05
0.80
0.98
0.99
0.99
0.20
0.63
0.55
0.70
0.64
0.46
0.95
0.95
0.95
a
4 2 1
Estim ation was based on experim ents conducted with an initial concentration of 0.63 m M MnO k .
-; k ) 102
FIGURE 4. Arrhenius plot for the oxidation of TCE (0.06 mM) by
potassium permanganate (1 mM) in a phosphate-buffered solution
at ionic strength 0.05 M and pH 7.1. Error bars are standard deviation
for each estimated second-order rate constant k .
FIGURE 3. Dependence of the pseudo-first-order rate constants
3a, k3b, and k3c) on pH for oxidation of carboxylic acids. Kinetic
data were obtained by reacting 0.09 mM TCE with 0.63 mM MnO
(k
-
4
1p
in a phosphate-buffered aqueous solution at ionic strength 0.05 M.
transfer interaction rather than net charge (24, 12). The
interaction is initiated by a “slipping” of the metal toward
one end of the alkene double bond leaving the other end
more electropositive to reaction with the oxygen in per-
manganate. The proposed slipping motion along the carbon-
carbon double bond in the transition state apparently favors
trans among dichloroethene (DCE) isomers but is relatively
difficult to be induced when ethene is fully substituted by
four chlorines (a strong electron-withdrawing group), such
as tetrachloroethene (PCE). This mechanism is supported
by our previous results that showed trans-DCE with much
greater reactivity than cis-DCE and PCE with a reaction rate
more than 1 order of magnitude smaller than TCE, cis-DCE,
and 1,1-DCE (7).
tion of an organometallic complex between the reactants
and the intermediate (14).
In the classical view, as summarized in ref 10, the
formation of cyclic ester involves a 3+2 electrocyclic addition
of permanganate ion to the carbon-carbon double bond
(
19)
where the superscript q indicates an activated complex at
transition state. The modern ideas on alkene oxidation by
permanganate have been advanced by refs 14 and 23. The
initial reaction proceeds with 2+2 insertion of an alkene
π-bond directly into a metal-oxo bond of manganese, forming
a transient organometallic complex (14) followed by rear-
rangement to the cyclic ester (12, 15). This modern theory
was confirmed by showing that relatively lower activation
energy was required for the 2+2 reaction pathway as
compared to 3+2 (23). With the modern theory, TCE is more
likely attacked by electropositive metal (manganese) via a
transition state:
q
The small enthalpy ∆H (39 kJ/ mol) and negative entropy
q
∆S (-14 J/ mol) are consistent with the observations of the
oxidation of other alkenic substrates by permanganate under
various conditions, even though their products are different.
The results suggest that TCE oxidation proceeds via similar
transition states but that these transformation products are
subjected to various post-transition-state reactions, which
are highly dependent on reaction conditions.
The existence of various reaction pathways is evidenced
by the formation of different intermediate products as a
function of pH. At pH 4, the decomposition of cyclic ester
2
involves both oxidative hydrolysis (followed by carbon-
carbon double bond cleavage) to form formic acid and
hydrolysis (with subsequent reactions) to form oxalic acid
(
20)
(
Figure 1). The calculated ratios of rate constants (k2a/ k
2
and
k
2c/ k
2
) indicate that the former reaction pathway is pre-
This reaction mechanism more reasonably explains why
permanganate ion with electron-rich oxygen termini could
attack the carbon-carbon double bond electrophilically. The
attack is controlled by orbital overlap via delocalized charge-
dominant (77% of total production) while the latter one is
much less important (20%). With increasing pH, almost all
the cyclic esters are hydrolyzed via 6. Products are further
decomposed via oxidative hydrolysis, hydrolysis, and electron
2
5 4 0
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 34, NO. 12, 2000

transfer to form acyl halides, which are very reactive and
easily hydrolyzed to glyoxylic and oxalic acids (95-97%).
Decomposition of acyclic ester 7 via 8 to form glycolic acid
high oxidation rate of carboxylic acids at low pH, as illustrated
in Figure 2, CO
pH.
2
is accumulated more rapidly with decreasing
2 1
is insignificant. The estimated k / k ratio from our kinetic
Acknowledgments
data also supports the postulate that the decomposition of
cyclic ester is much faster than its formation. The series of
reactions, during the fast ester decomposition, liberate almost
all the chlorine substitutes on TCE (7).
In the final step, the oxidation of carboxylic acids is
relatively slow as compared to the previous steps. The
oxidation may involve hydride abstraction by permanganate
ion (25) and carbon-carbon bond cleavage. In the aqueous
solution, carboxylic acids can be dissociated:
We thank Drs. O. H. Tuovinen, Y. P. Chin, and S. Traina for
their advice and assistance. This material is based upon work
supported by the U.S. Department of Energy under Grant
DE-FG07-96ER14735.
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(
(
(
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(
23)
(
7) Yan, Y. E.; Schwartz, F. W. J. Contam. Hydrol. 1999, 37, 343-
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(
(9) Wiberg, K. B.; Saegebarth, K. A. J. Am. Chem. Soc. 1957, 79,
2
822-2824.
(
(
(
10) Freeman, F. Rev. React. Species Chem. React. 1976, 1, 179-226.
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a b
where K , K , Kc1, and Kc2 are ionization constants for formic,
glyoxylic, and oxalic acids, respectively. If all species are
considered to react independently with permanganate, eqs
(13) Wolfe, S.; Ingold, C. F. J. Am. Chem. Soc. 1981, 103, 938-939.
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7
-9 can be further developed to eqs 25-27 (Table 1). As
1
977, 99, 3120-3128.
15) Freeman, F.; Kappos, J. C. J. Am. Chem. Soc. 1985, 107, 6628-
633.
compared to eqs 25-27 with eqs 7-9, the second-order rate
constants for oxidation of formic, glyoxylic, and oxalic acids
can be simplified as
(
(
6
16) Kivinen, A. In The Chemistry of Acyl Halides; Patai, S., Ed.;
Wiley: New York, 1972; pp 177-230.
(
(
17) Berka, A.; Koreckova, J. Anal. Lett. 1973, 6, 1113-1123.
18) Rodriguez, J.; Sanchez Burgos, F. Ion (Madrid) 1975, 35, 241-
245.
k3ap1
+
k3ap
)
)
[H ] + k
(28)
(29)
3ap2
3bp2
Ka
(
19) Mahmood, A. J.; Begum, M. Dacca Univ. Stud., Part B 1975, 23,
5
1-64.
k3bp1
(20) Huang, K. C.; Hoag, G. E.; Chheda, P.; Woody, B. A.; Dobbs, G.
M. Environ. Eng. Sci. 1999, 16, 265-274.
(21) Perez-Benito, J. F.; Arias, C.; Brillas, E. Int. J. Chem. Kinet. 1990,
+
k3bp
[H ] + k
Kb
2
2, 261-287.
(
22) Wagner, G. J. Russ. Phys. Chem. Soc. 1895, 27, 219-236.
+
2
+
k3cp1[H ] + k
K[H ] + k
K K
3cp3 cl c2
(23) Rappe, A. K.; Goddard, W. A., III. J. Am. Chem. Soc. 1982, 104,
448-456.
(24) Toyoshima, K.; Okuyama, T.; Fueno, T. J. Org. Chem. 1980, 45,
3cp2
k3cp
)
(30)
+
K K + K [H ]
cl c2
c1
1
600-1604.
(
25) Gardner, K. A.; Mayer, J. M. Science 1995, 269, 1849-1851.
when pH is much higher than pK
a
, pK
b
, and pKc1, resulting
+
+
+ 2
+
in [H ] , K
a
, [H ] , K
b
, and [H ] , Kc1
Kc2 and Kc1[H] . The
Received for review November 15, 1999. Revised manuscript
received March 20, 2000. Accepted March 27, 2000.
relationship between logarithms of second-order rate con-
stants with pH shown in Figure 3 is consistent with the
mechanism described by eqs 28-30. Because of the relatively
ES991279Q
VOL. 34, NO. 12, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
2 5 4 1