Kinetics and mechanism of the reduction of sodium chlorite by sodium hydrogen ascorbate in aqueous solution at near neutral pH

Article abstract of DOI:10.1002/kin.20847

The reduction of chlorite ion by the hydrogen ascorbate ion in a neutral solution safely produces chlorine dioxide. The decrease in absorbance at 268 nm with the presence of dimethyl sulfoxide (DMSO) allows measurement of the ascorbate disappearance in the reaction with excess chlorite. The measured rate constant at 25 ± 0.02°C, 3.67 × 10^-4 M DMSO, ionic strength 0.51 M (NaClO₄), and in the presence of 3.32 × 10 ^-9 M ethylenediaminetetraacetic acid is 13.81 ± 1.30 M ^-1 s^-1. Rate constant measurements over the range 15-35°C gave an Arrhenius activation energy of 75.51 ± 4.53 kJ mol^-1. This result is the first reported determination of the kinetics of this reaction and is consistent with either electron- or oxygen-transfer mechanisms. Anomalously, reduction of chlorite results in its oxidation, because intermediate hypochlorite oxidizes chlorite.

Full text of DOI:10.1002/kin.20847

Kinetics and Mechanism
of the Reduction of Sodium
Chlorite by Sodium
Hydrogen Ascorbate
in Aqueous Solution at
Near Neutral pH
MARIA CURTIN,1 SEAN DWYER,1 DINO BUKVIC,1 CHRISTOPHER J. DOONA,2 KENNETH KUSTIN3
1Department of Chemistry, Stonehill College, Easton, MA 02357
2U.S. Army-Natick Soldier RDEC, Natick, MA 01760
3Department of Chemistry, Emeritus, Brandeis University, Waltham, MA 02453
Received 1 July 2013; revised 11 January 2014
DOI 10.1002/kin.20847
Published online 3 February 2014 in Wiley Online Library (wileyonlinelibrary.com).
ABSTRACT: The reduction of chlorite ion by the hydrogen ascorbate ion in a neutral solution
safely produces chlorine dioxide. The decrease in absorbance at 268 nm with the presence of
dimethyl sulfoxide (DMSO) allows measurement of the ascorbate disappearance in the reaction
with excess chlorite. The measured rate constant at 25 0.02◦C, 3.67 × 10−4 M DMSO, ionic
strength 0.51 M (NaClO4), and in the presence of 3.32 × 10−9 M ethylenediaminetetraacetic acid
is 13.81 1.30 M−1 s−1. Rate constant measurements over the range 15–35◦C gave an Arrhenius
activation energy of 75.51 4.53 kJ mol−1. This result is the first reported determination of the
kinetics of this reaction and is consistent with either electron- or oxygen-transfer mechanisms.
Anomalously, reduction of chlorite results in its oxidation, because intermediate hypochlorite
ꢀC
oxidizes chlorite. 2014 Wiley Periodicals, Inc. Int J Chem Kinet 46: 216–219, 2014
few seconds. How does the addition of a reductant to a
INTRODUCTION
strong oxidant oxidize the oxidant? In this case, the an-
swer is through formation of the reactive intermediate
hypochlorite [ClO−; Cl(I)], which once formed oxi-
dizes chlorite to chlorine dioxide autocatalytically [1].
The kinetics of the initial reaction is difficult to follow,
because hypochlorite, in addition to reacting with chlo-
rite, also reacts with and oxidizes all carbon-containing
species in solution. Dimethyl sulfoxide (DMSO),
The addition of the reductant ascorbate (AH−) to ex-
cess aqueous chlorite [ClO−2 ; Cl(III)] produces copious
amounts of chlorine dioxide [ClO2; Cl(IV)] within a
Correspondence to: Maria Curtin; e-mail: mcurtin@
stonehill.edu.
C
ꢀ
2014 Wiley Periodicals, Inc.

KINETICS AND MECHANISM OF THE REDUCTION OF NaClO2 BY NaAH
217
2 × 10−9 M EDTA. Separately, 0.10215 g of NaAH
was dissolved and diluted with water in a 100.0-mL
volumetric flask to produce a 5.16 × 10−4 M solution.
For isothermal kinetic studies at different constant
temperatures, solutions were mixed manually in a cu-
vette placed in the thermostated holder of the UV–vis
spectrophotometer. A 3.00-mL aliquot of the NaClO2
solution was transferred to the cell and allowed to ther-
mally equilibrate for about 10 min with vigorous stir-
ring and a constant flow of argon gas passing through
the sample compartment to reduce available oxygen.
An 80-μL aliquot of DMSO was added to the NaClO2
solution in the cuvette, followed by 20 μL of NaAH so-
lution, and the absorbance versus time data collection
was started. The temperature range for these experi-
ments was held between T = 15 and 35◦C. At temper-
atures below this range, condensation beclouded the
cuvette, and at temperatures above this range, the reac-
tion was too fast at these concentrations to accurately
measure the kinetics. In some experimental conditions
to reduce the NaClO2 concentration, the volume of
NaClO2 solution added was reduced, accordingly, and
the volume and ionic strength of the final solution in
the cuvette were kept constant by correspondingly in-
creasing the amount of 0.500 M NaClO4 solution prior
to the addition of DMSO and NaAH.
however, scavenges hypochlorite very efficiently [2]
and allows spectrophotometric determination of the ki-
netics of the reduction of chlorite by ascorbate, which
is reported herein.
EXPERIMENTAL
A Cary UV–vis spectrophotometer with a Peltier ther-
mostated ( 1◦C) and a stirred cell was used to monitor
the reaction progress at 268 nm. The experimentally
measured molar absorptivities at 268 nm for sodium
ascorbate and sodium chlorite in deionized water, 1.2
× 104 and 1.3 × 102 M−1 cm−1, respectively, show
that in the presence of excess chlorite the decrease in
the absorbance is due to the disappearance of ascor-
bate. Eppendorf Research pipettes (1−20 μL and 1–5
mL volumes) were used to add reagents.
Ethylenediaminetetraacetic acid (EDTA; used to
complex any trace metals present in solution) and
DMSO were obtained from Fisher Scientific, Pitts-
burgh, PA. Sodium perchlorate (NaClO4; used to main-
tain a constant ionic strength, μ) and sodium hydrogen
ascorbate (NaAH) were obtained from Acros Organ-
ics, distributed by Fisher Scientific, and sodium chlo-
rite (NaClO2; 80% technical grade) was purchased
from Labchem, Inc., Zelienophe, PA. Type I water
(18.2 Mꢀ) produced with a Waters Milli-Q A-10 water
filtration system was used to make all solutions.
RESULTS AND DISCUSSION
With the exception of NaClO2, all reagents were
used without further purification. The NaClO2 solid
was dissolved in an equal amount of water to make a
50:50 w/w solution. The mixture was heated slowly to
T = 50◦C and held at that temperature until most of
the NaClO2 dissolved. The concentrated NaClO2 solu-
tion was filtered using a fritted glass filter of medium
porosity. The clear solution was allowed to cool and
crystallize for several days. The solid formed was fil-
tered and dried under vacuum to remove any remaining
water. About 10% of the initial solid was obtained and
stored in a desiccator (Drierite). The purified NaClO2
was analyzed using a standard iodometric technique [3]
and found to be 97% pure.
For kinetic studies, separate stock solutions of
NaClO2 and NaAH were typically prepared as follows,
with minor variations based on appropriate changes in
the starting concentrations: Solutions of NaAH were
prepared fresh daily, and NaClO2 solutions were pre-
pared every third day. For a typical run, NaClO2
(0.12514 g) and NaClO4 (30.620 g) were dissolved in
water. One milliliter of 1.0 × 10−6 M EDTA stock solu-
tion was added, and the mixture was diluted with water
to 500.0 mL in a volumetric flask to make a solution
containing 0.00276 M NaClO2, 0.500 M NaClO4, and
In the presence of DMSO and at near-neutral pH (6–
8), the overall reduction of ClO−2 by AH− proceeds
according to the stoichiometry of reaction (I),
ClO−2 + 2AH− = C1− + 2A + 2OH−
(I)
where A is dehydroascorbic acid. Each absorbance ver-
sus time profile of reaction (I) displayed pseudo–first-
order kinetics, where kobs is the pseudo–first-order rate
constant and k2 = kobs/[ClO−2 ]. Holding all other re-
action conditions constant, the effect of varying the
amount (in μL) of DMSO added was studied (see
Table I), and 80 μL DMSO was chosen to continue
the kinetic studies. To correct observed rate constants
for any systematic error due to incomplete initial mix-
ing, a double-reciprocal plot of observed rate constant
(1/kobs, s) versus concentration of excess ClO2−
(1/[ClO−2 ], M−1) was carried out [4]. As shown in Fig.
1, the reciprocal of the slope of the plot yields the value
for the second-order rate constant k2 = 13.81 2%
M
−1 s−1 at T = 25◦C. The rate constants determined at
different temperatures were used to obtain values for
the Arrhenius parameters. Specifically, a plot of kobs
versus 1/T (K−1) (see Fig. 2) yielded values for the
International Journal of Chemical Kinetics DOI 10.1002/kin.20847

218
CURTIN ET AL.
imparted through the presence of the {AH2 − AH− −
Table I DMSO-Dependence of Observed Rate
Constanta
A
2−} conjugate acid–base pairs, which tended to keep
pH near neutral.
μL of DMSO
added
DMSO
(mol/L)
Kobs (s−1
)
Kobs avg
Interestingly, there was an earlier study that es-
timated the value of k2 [5]. Pulse radiolysis experi-
ments were performed by Huie and Neta to measure
electron-transfer reactions of ClO2 and NO2. Their at-
tempts to produce ClO2 by OH oxidation of ClO2−
in the presence of AH− were thwarted, because of a
thermal reaction between ClO−2 and AH−. Based on
this observation, they inferred a rate constant of k2 =
34 M−1 s−1, which is not very different from the mea-
sured value reported here. The importance of their
study was to distinguish, for example, the autoxida-
tion of AH−, which is very slow [6], from the AH−
oxidation by ClO−2 , which is rapid. The present Arrhe-
nius parameters, moreover, indicate that the first step
in the overall reaction proceeds with a high collision
frequency and modest Ea.
40
80
100
120
180
0.18
0.36
0.45
0.53
0.79
0.0225
0.0223
0.0262
0.0264
0.0263
0.0247 ( 7.6%)
aReaction conditions of μ = 0.5 M (NaClO4), [EDTA] =
3.33 × 10−9 M, [AH−] = 3.33 × 10−5 M, [ClO−2 ] = 3.33 × 10 −3 M,
and T = 25 1◦C.
For many one-electron acceptors, the mechanism
of AH− reduction typically proceeds via outer-sphere
electron-transfer [7]. A valid mechanism for ClO−2 oxi-
dation with many other organic reductants, however, is
adduct formation and oxygen-atom transfer [8]. Both
mechanisms may plausibly be occurring in the (ClO−2 –
AH−) reaction.
The possible existence of electron transfer in the
(ClO−2 –AH−) reaction can be tested by applying the
Marcus cross-reaction theory to a plausible first step
in a reduction mechanism [9]. The cross-reaction of
interest is the redox reaction (R12):
Figure 1 Double reciprocal plot, kobs versus [ClO−2 ]: o,
experiment; —, regression line, 1/kobs = 17.84 + 0.072/
[ClO−2 ].
AH− + ClO−2 = A−• + ClO + OH−, k12 (R12)
where AH− is the ascorbate ion, A−• is the ascorbyl
radical ion, and k12 = k2. The forward (cross) rate
constant, k12, is related to the rate constants, k11 and
k22, of the following two electron exchange reactions
(E1) and (E2):
AH− + A− = A− + AH−, k11
ClO−2 + ClO = ClO2− + ClO, k22
(E1)
(E2)
Figure 2 Arrhenius plot, ln kobs versus 1/T: •, experiment;
—, regression line, ln kobs = 26.53–9082 (1/T).
Cross-reaction (R12) simplifies the actual situation,
as exemplified by pulse-radiolysis studies of oxoan-
ions [10,11]. In these systems, the initial step in the
reaction between the hydrated electron and the oxoan-
ion involves formation of an electron adduct; namely
ClO=2 , that dissociates at a time scale on the or-
der of nanoseconds to produce ClO and hydroxide
ion (OH−). The self-exchange equation (E1) recog-
nizes that the ascorbyl free radical is a strong acid
Arrhenius parameters A = 3.32 × 1011 M−1 s−1 and
activation energy Ea = 75.51 kJ mol−1) with relative
error of 6%.
The attempts to determine the pH dependence of kobs
using common buffer systems to control pH at different
levels tended to be compromised by buffer interfering
with components of the reaction mixture. In general
for the studies reported here, some pH buffering was
International Journal of Chemical Kinetics DOI 10.1002/kin.20847

KINETICS AND MECHANISM OF THE REDUCTION OF NaClO2 BY NaAH
219
(pKa = 0.45) [12], and at pH 7 would be fully dis-
sociated. The self-exchange equation (E1) omits the
hydroxide ion.
The full Marcus cross-reaction rate constant equa-
tion contains Coulombic and non-Coulombic work
terms, which can be neglected for the set of reactions
presently being considered. Accordingly, the follow-
ing two equations describe the relationships between
the cross-reaction and the self-exchange reaction rate
constants:
agogic as well as practical benefits, including impor-
tant ramifications for production of chlorine dioxide
used to decontaminate medical instruments, fresh pro-
duce, and food contact surfaces in austere environ-
ments. Chlorine dioxide is also one of the preferred
oxidants used in wastewater treatment. It is an effec-
tive oxidant of harmful organic contaminants [15,16].
This new method of producing chlorine dioxide using
ascorbate and chlorite is a more environmentally sus-
tainable and a safer source of ClO2 when compared
to current methods that produce ClO2 by either the re-
action of chlorite with chlorate or the acidification of
chlorite with a strong acid such as sulfuric acid [17].
ꢀ
k12
=
k11k22K12f
(M1)
(M2)
lnf = (lnK12)2/4ln(k11k22/Z2)
The authors gratefully acknowledge the SURE program at
Stonehill College for support of Sean Dwyer and NSF-STEP
grant #DUE-0622540 and the Stonehill College Department
of Chemistry for support of Dino Bukvic.
in which K12 is the equilibrium constant for reaction
(R12) and Z is the collision frequency with a value
equal to 1 × 1011 M−1
s
−1. The purpose of this test is
to find the value of k22 and determine whether its value
is plausible. To carry out this analysis, it is necessary
to know the value of K12, which is considered next.
Because ClO is unstable in aqueous solution, it is
difficult to devise a half-cell reaction from which an
appropriate E0 value can be measured. One successful
kinetic method measures the rate of equilibration for
the reaction between the pulse radiolytically generated
carbonate radical ion and hypochlorite ion [13]. Taking
this kinetic-determined value of 1.41 V forꢁthe couple
ClO/ClO− leads to a calculated value of E0 = 0.69 V
for the ClO−2 /ClO couple at pH 7 with respect to the
normal hydrogen electrode. The calculated value for
the equilibrium constant K12 ( = 1.18 × 105) following
from this determination of E0 is used in the calculation
of k22.
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International Journal of Chemical Kinetics DOI 10.1002/kin.20847