Chlorine dioxide-facilitated oxidation of the Azo dye amaranth

Article abstract of DOI:10.1021/jp206175s

The oxidation reaction of amaranth (trisodium 2-hydroxy-1-(4-sulfonato-1- naphthylazo)naphthalene-3,6-disulfonate or AM^-) by chlorine dioxide (ClO₂) in aqueous conditions was investigated in detail. The major reaction products immediately after decolorization of AM^- were 1,2-naphthoquinone disulfonate sodium salt and 1,4-napthalenedione. The reaction had first-order dependence on both AM^- and ClO₂. The rate-limiting step involved the reaction between AM^- and OH ^- ions. The role of hydroxide ion as a catalyst was established. The second-order rate constant increased with pH, from (19.8 ± 0.9) M ^-1 s^-1 at pH 7.0, (97.1 ± 2.3) M^-1 s^-1 at pH 8.0 to (132.5 ± 2.8) M^-1 s^-1 at pH 9.0. In the pH range of 6.0-7.5, the catalytic constant for OH^- ion was 4.0 × 10⁹ M^-2 s^-1. The energy and entropy of activation values for the reaction were 50.0 kJ mol^-1 and -658.7 J K^-1 mol^-1, respectively. A probable reaction mechanism was elucidated and was validated by simulations.

Full text of DOI:10.1021/jp206175s

ARTICLE
pubs.acs.org/JPCA
Chlorine Dioxide-Facilitated Oxidation of the Azo Dye Amaranth
S. Nadupalli, N. Koorbanally, and S. B. Jonnalagadda*
School of Chemistry, University of KwaZulu-Natal, Westville Campus, Chiltern Hills P. Bag 54001, Durban-4000, South Africa
S Supporting Information
b
ABSTRACT: The oxidation reaction of amaranth (trisodium
2-hydroxy-1-(4-sulfonato-1-naphthylazo)naphthalene-3,6-dis-
ulfonate or AMꢀ) by chlorine dioxide (ClO2) in aqueous
conditions was investigated in detail. The major reaction
products immediately after decolorization of AMꢀ were 1,2-
naphthoquinone disulfonate sodium salt and 1,4-napthalene-
dione. The reaction had first-order dependence on both AMꢀ
and ClO2. The rate-limiting step involved the reaction between AMꢀ and OHꢀ ions. The role of hydroxide ion as a catalyst was
established. The second-order rate constant increased with pH, from (19.8 ( 0.9) Mꢀ1 sꢀ1 at pH 7.0, (97.1 ( 2.3) Mꢀ1 sꢀ1 at pH
8.0 to (132.5 ( 2.8) Mꢀ1 sꢀ1 at pH 9.0. In the pH range of 6.0ꢀ7.5, the catalytic constant for OHꢀ ion was 4.0 ꢁ 109 Mꢀ2 sꢀ1. The
energy and entropy of activation values for the reaction were 50.0 kJ molꢀ1 and ꢀ658.7 J Kꢀ1 molꢀ1, respectively. A probable
reaction mechanism was elucidated and was validated by simulations.
’ INTRODUCTION
with prolonged exposure, unless other advanced oxidation pro-
cesses are employed.20
Chlorine dioxide is commonly used as a bleaching agent in the
paper and pulp and other industries.1ꢀ4 It is also employed in
water treatment for oxidative degradation of organic and inor-
ganic pollutants.5,6 Synthetic and natural dyes are in great demand in
the chemical and other manufacturing industries. Many of these
dyes are toxic and suspected carcinogens and mutagens, and pose
environmental problems when they enter water systems.7,8 The
azo dyes are reported to form about 60ꢀ70% of the dyes used in
the textile industry.9 The textile industry generates a large quantity
of wastewater containing organic matter, suspended solids,
metals, halogenated organic surfactants and other pollutants.10,11
Thus, the fate of these dyes in water systems and their elimina-
tion or removal poses new challenges in the wastewater treat-
ment. For color removal from dye material containing waste-
waters, in addition to the use of adsorption, several technologies
such as chemical oxidative degradation, catalyzed oxidation, elec-
trochemical treatment10,12,13 and bio- and photodegradation14,15
were reported. A chlorine dioxide initiated microwave-enhanced
oxidation process was also reported for the oxidation of azo
dyes;14 however, not much is known about the reactivity of ClO2
under various pH ranges. The extent of dye degradation always
depended on the technique adopted and the duration of ex-
posure. Some partially degraded products could be hazardous on
release into the ecosystem.15,16 Chlorine dioxide exists as a free-
radical monomer.17 The chlorineꢀoxygen bonds have a double-
bond character, an OꢀClꢀO angle of about 117.5°, and a
chlorineꢀoxygen bond length of 0.147 nm.18 Chlorine dioxide
typically reacts with organics as an oxidant with little or no
chlorination. The breaking of organic carbonꢀcarbon bonds is
generally not extensive in most of the reactions, so that the
oxidation byproducts are also organics.19 Complete mineraliza-
tion is not achievable with chlorine dioxide as the oxidant even
Amaranth (AMꢀ), which is used in the textile industry, is a sus-
pected carcinogen and is reported to have an LD50 of 1000 mg kgꢀ1
in rat and mouse models in material safety data sheets (MSDS).
It is still used widely in the food and dye industry, and there is no
reported toxicity to humans to date. It is commonly referred to as
Red No. 2, Food Red 9, Acid Red 27, and Azorubin S. It is a
water-soluble azo dye, with a sharp absorption peak in the visible
region (λmax = 620 nm, ε = 2.1 ꢁ 104 dm3 molꢀ1 cmꢀ1) with no
peak shift due to pH change.21 This communication focuses on
the mechanism of oxidative decolorization of AMꢀ by chlorine
dioxide under aqueous conditions by kinetic approach.
’ EXPERIMENTAL SECTION
Amaranth (Aldrich), sodium hydroxide (BDH), 98% sulfuric
acid (BDH), and anhydrous sodium sulfate (Univar) were used
as supplied. All experiments were performed using double
distilled water and analytical-grade or high-purity reagents. The
chlorine dioxide was generated by standard procedure by acid-
ifying the aqueous solution of sodium chlorite.22 The stock solu-
tion of ClO2 was stored in the dark, at room temperature, and
was standardized before the reaction, by idometric titration with
standard sodium thiosulfate solution.
Kinetic Measurements. All the kinetic measurements were
conducted monitoring the depletion of amaranth by using the
HI-TECH micro volume double mixing stopped-flow apparatus.
The kinetic data acquired at single wavelength was analyzed
using the KinetAsyst 3.10 software, which allows the matching of
Received: June 30, 2011
Revised:
September 14, 2011
Published: September 18, 2011
r
2011 American Chemical Society
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Figure 1. Absorbance versus time plot for the reaction of [AMꢀ]0 (7.0 ꢁ 10ꢀ5 M) with [ClO2]t (1.15 ꢁ 10ꢀ3 M) at pH = 9.0 at λmax = 620 nm.
Table 1. Reaction between Amaranth and Chlorine Dioxide
at Constant Ionic Strength
Table 2. Calculated [OHꢀ]eq Values and Corresponding
Second Order Constants for the Reaction of [AM]0 (7.0 ꢁ
10ꢀ5 M) with [OHꢀ]eq/M (7.94 ꢁ 10ꢀ9 to 7.94 ꢁ 10ꢀ6
)
[ClO2]/10ꢀ3 M
k0/sꢀ1a
[OHꢀ]eq/M
k0/sꢀ1a
k/102
M
ꢀ1 sꢀ1
2.52
3.79
5.05
6.31
7.57
3.320
3.740
3.990
4.340
4.800
7.94 ꢁ 10ꢀ9
1.58 ꢁ 10ꢀ8
2.51 ꢁ 10ꢀ8
6.31 ꢁ 10ꢀ8
1.78 ꢁ 10ꢀ7
6.31 ꢁ 10ꢀ7
3.16 ꢁ 10ꢀ6
7.94 ꢁ 10ꢀ6
0.08
0.12
0.15
0.17
0.49
1.30
4.30
5.80
0.3
0.5
1.0
1.8
3.3
6.3
a Mean of four replicate experiments with relative standard deviation <4%.
experimental results with different rate equations and to estimate the
rate constants by choosing appropriate integrated rate equations.23
10.4
12.9
a Mean of four replicate experiments with relative standard deviation <4%.
’ RESULTS AND DISCUSSION
Chlorine dioxide is relatively inert under acidic conditions, but
under neutral and basic pH it becomes more reactive.25 Reac-
tions involving chlorine dioxide as an oxidant proceed through
free radical electrophilic abstraction, rather than by oxidative sub-
stitution or addition, which happens with chlorine or hypochlorite.26
Due to its exceptional one-electron abstraction ability, chlorine
dioxide acts as a highly selective oxidant and reduces to chlorite
(ClO2ꢀ). To assess the role of pH on the reaction order with
respect to chlorine dioxide, a series of experiments were conducted
with varied initial concentrations of ClO2 at pH 7.0, 8.0, and 9.0.
The values of pseudo-first-order rate coefficients, k0, for different
[ClO2]0 at different [OHꢀ]0 values are summarized in Table 2.
The linear plots of ln k0 versus ln[ClO2]0 at fixed pH 7.0, 8.0, and
9.0 gave slopes equal to 1.08, 1.03, and 1.01, respectively (Figure 2),
suggesting that the change in pH does not influence the order with
respect to oxidant or the overall reaction mechanism.
Although, a number of literature reports describe an increased
decomposition of chlorine dioxide at higher pH conditions,27ꢀ29
very few studies explain the chemistry behind its increased oxidative
activity. The step involving the one electron reduction of chlorine
dioxide forming chlorite ion is fast (ClO2(aq) + e‑ f ClO2ꢀ, E° =
0.954 V).30 At alkaline pH, even in the absence of reducing
substrates, chlorine dioxide is known to disproportionate to chlorite
and chlorate ions. Under alkaline conditions, the reactions that
facilitate the chlorine dioxide disproportionation only lead to the
formation of less reactive species such as chlorite and chlorate.31
Reaction Kinetics. (a). Order with Respect to Amaranth and
Chlorine Dioxide. The kinetics of the reaction between amaranth
and chlorine dioxide was studied under varied conditions with
large excess concentrations of chlorine dioxide and other
reagents, relative to the fixed initial concentration of dye. All the
experiments were carried out at 25 °C ((0.1) except the
experiments for estimation of the energy parameters. Figure 1
shows a typical curve with [ClO2]0 = 1.15 ꢁ 10ꢀ2 M and [dye]0
= 7.0 ꢁ 10ꢀ5 M and pH 9.0. Perusal of Figure 1 shows that the
depletion of amaranth followed exponential decay, and the
reaction was completed in less than two seconds, suggesting that
the reaction was very fast. Employing the kinetic fit software and
using the first-order equation y = ꢀA exp(ꢀkx) + C, the series of
curves analyzed show that the test-fit results of the curve is a fair
agreement with small residuals, confirming that the reaction
order with respect to the dye is one. Table 1 summarizes the values
of the pseudo-first-order rate coefficient k0 obtained for different
initial concentrations of chlorine dioxide at fixed ionic strength.
The plot of ln[ClO2] versus ln k0 gave a linear curve with a slope
of 1.07 and correction coefficient of 0.99, suggesting that the
reaction rate has first-order dependence on [ClO2]0.
(b). Effect of pH on Reaction Rate and Reaction Order with
Respect to Hydroxide Ion. The literature survey shows that,
unlike Cl2 and HOCl, which are more effective oxidants under acidic
conditions, chlorine dioxide is more reactive at alkaline pH.24
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ARTICLE
Table 3. Observed Rate Constants at pH 7.0, 8.0, and 9.0 for
the Reaction of [AM]0 (7.0 ꢁ 10ꢀ5 M) with [ClO2]t (2.52 ꢁ
10ꢀ3 to 3.53 ꢁ 10ꢀ3 M)a
pH = 7.0
[ClO2] ꢁ 10ꢀ3 /M
k0/ sꢀ1b
k2/Mꢀ1sꢀ1
k
2.52
2.77
3.03
3.28
3.53
0.39
15.77
16.06
1.58 ꢁ 108
1.61 ꢁ 108
1.68 ꢁ 108
1.64 ꢁ 108
1.61 ꢁ 108
0.44
0.50
16.76
0.53
0.57
16.42
16.14
mean k2 with std.dev.
16.2 ( 0.37
Figure 2. Plot of ln k0 versus ln[ClO2] for the reaction of [AMꢀ]0
(7.0 ꢁ 10ꢀ5 M) with [ClO2] (0.025ꢀ0.035 ꢁ 10ꢀ3 M): (a) pH = 9.0,
(b) pH = 8.0, (c) pH = 7.0.
pH = 8.0
[ClO2] ꢁ 10ꢀ3 /M
k0 /sꢀ1
k2/Mꢀ1sꢀ1
k
2.52
2.77
3.03
3.28
3.53
2.50
99.20
93.86
9.9 ꢁ 107
9.4 ꢁ 107
9.6 ꢁ 107
9.8 ꢁ 107
9.9 ꢁ 107
2.60
ClO2 þ AMꢀ f ClO2 þ AM
ð1Þ
ð2Þ
ꢀ
2.90
95.70
3.20
3.50
97.56
2ClO2 þ 2OHꢀ f ClO2 þ ClO3 þ H2O
ꢀ
ꢀ
99.15
mean k2 with std.dev.
97.1 ( 2.3
Under alkaline conditions, chlorite and chlorate ions are less re-
active relative to chlorine dioxide, and chlorate ion is almost inert.
The reactions of chlorite are slow under acidic conditions and
much slower at alkaline pH. Thus, the other reactive species that
could attack the substrates at a significant rate are possibly
HClO2, Cl2O4, HOCl, and peroxide ion. Their concentrations
under alkaline pH determine their contribution toward higher
reactivity. The concentrations of HClO2 and HOCl under alka-
line conditions will be very small, thus contributing negligibly
toward the oxidation of the substrate. The formation of Cl2O4, a
dimer of ClO2, is not pH dependent. Hydroxide ion with two
molecules of ClO2 could lead to the formation of peroxide ion,
but in small concentrations.32,33 The reaction chemistry under
alkaline conditions can be expressed in the following equations:
pH = 9.0
[ClO2] ꢁ 10ꢀ3 /M
k0 /sꢀ1
k2/Mꢀ1sꢀ1
k
2.52
2.78
3.03
3.28
3.53
3.32
131.74
134.53
1.3 ꢁ 107
1.3 ꢁ 107
1.3 ꢁ 107
1.3 ꢁ 107
1.4 ꢁ 107
3.74
3.99
131.68
4.22
4.80
128.65
135.97
mean k2 with std.dev.
132.5 ( 2.8
a k2 = k0/[ClO2]; k = k0/{[ClO2][OHꢀ]}. b Mean of four replicate
experiments with relative standard deviation (RSD) < 4%.
̅
ClO2 þ OHꢀ h ðOClOOHÞꢀ
ðOClOOHÞꢀ þ ClO2 h OClOOH þ ClO2
OClOOH þ OHꢀ h HOClO þ HOOꢀ
HOOꢀ þ 2ClO2 þ OHꢀ h 2ClO2 þ H2O
ð3Þ
ð4Þ
ð5Þ
ð6Þ
ꢀ
ꢀ
Therefore, the increased reactivity of ClO2 under alkaline pH
certainly could not be due to such less reactive species. Hence,
the enhanced reactivity observed at high pH needs to be explained
differently. This could happen through the generation of many
active species by the OHꢀ ion. Possibly, for such species, a transient
complex involves OHꢀ ion, ClO2, and substrate. Considering the
first-order dependence of the reaction rate on hydroxide ion near
neutral conditions, its direct role in the rate-limiting step can be
envisaged, through the probable formation of a reactive inter-
mediate involving oxidant, organic substrate, and hydroxide ion.
Figure 3. Plot of log k0 versus log [OHꢀ] at different pH conditions.
summarized in Table 3. A perusal of ln k0 versus ln[OHꢀ] plots in
Figure 3 shows that with increasing pH, the reaction order with
respect to [OHꢀ] decreases. While the reaction order with
respect to hydroxide ion under near-neutral conditions was near
unity, it decreased with rising pH. This can be explained in terms
of two competitive oxidative reactions involving ClO2, but with
and without hydroxide ion. The reaction rate contributed by the
two competitive reactions may be expressed as
2ꢀ
ClO2 þ AMꢀ þ OHꢀ f ½ClO2AMOHꢂ
ð7Þ
To establish the role of hydroxide ion in the reaction mechanism,
under an otherwise similar environment, the kinetics of the
reaction was investigated at different pH conditions. The pseu-
do-first-order rate constants obtained at different pH ranges are
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Activation Parameters. The enthalpy and entropy of activa-
tion of a chemical reaction provide valuable information about the
nature of the transition state, and hence about the reaction mechan-
ism. The temperature dependence of the reaction rate constant
was estimated by performing the kinetic runs at temperature range
10ꢀ30 °C. The energy of activation and enthalpy of activatiꢀo1n
values for the reaction were 50.0 kJ molꢀ1 and 47.6 kJ mol
,
.
respectively, and the entropy of activation was ꢀ658.7 J Kꢀ1 molꢀ1
High negative entropy of activation suggests that transition state
requires the reacting molecules to orient into compact conforma-
tions and approach each other at a precise orientation.
Product Identification and Characterization. The crude
product (0.54 g) of the AMꢀClO2 reaction was separated using
column chromatography with silica gel as the stationary phase.
The mobile phase consisted of a hexane:dichloromethane step
gradient: 100% hexane (fractions 1ꢀ20), 10% dichloromethane
in hexane (fractions 20ꢀ30), and 30% dichloromethane in hexane
(fractions 30ꢀ40). Fractions of 10 mL were collected in each
step. The purification of fractions 30ꢀ40 gave a pure compound,
which showed the MS m/z with relative intensities of 318 (25%),
316 (100%), 300 (18%), 288 (28%), 282 (70%), 202 (10%), and
122(16%); product P1 was identified to be 1,2-naphthoquinone
disulfonate. The mass spec peak patterns agreed well with com-
pound obtained as the oxidation product of similar azo dyes with
hydrogen peroxide.36 Further purification of fractions 20ꢀ30 by
elution with 40% dichloromethane in hexane afforded a second
major compound, P2 (5 mg). The oxidation products were
identified as P1 = 1,2-naphthoquinone disulfonate sodium salt
and P2 = 1,4 naphthalenedione. (See Supporting Information
Figures S1, S2, and S3). While the oxidation product P1 is re-
ported to be nontoxic, P2 (1,4- naphthalenedione) is reported to
be highly toxic in nature compared to starting material with ORL-
Figure 4. Plot of log k0/[ClO2] versus log [OHꢀ] for the reaction
[AMꢀ]0 (7.0 ꢁ 10ꢀ5 M) with [ClO2]t (1.15 ꢁ 10ꢀ3 M), [OHꢀ]eq (1.0
ꢁ 10ꢀ8 to 6.31 ꢁ 10ꢀ7 M).
ꢀ
r ¼ k1½ClO2ꢂ½AMꢂ þ kOH ½ClO2ꢂ½OHꢀꢂ½AMꢀꢂ
ð8Þ
¼ fk1 þ kOH ½OHꢀꢂg½ClO2ꢂ½AMꢀꢂ
ꢀ
¼ k½ClO2ꢂ½AMꢀꢂ ¼ k0½AMꢀꢂ
ð9Þ
where k0 is the observed pseudo-first-order rate constant in the
presence of excess concentration of chlorine dioxide. The second-
orderrate constant, k, is equal to k0/[ClO2], and for fixed [ClO2]0
it can be expressed as k = k0/[ClO2] = {k1 + kOH [OHꢀ]},
ꢀ
where k1 is the second-order rate constant ꢀfor the reaction be-
tween chlorine dioxide and dye, while kOH is the third-order
rate constant for the OHꢀ ion catalyzed reaction between chlorine
dioxide and dye. Margerum et al. during their studies on the
oxidation of nitrogen dioxide with chlorine dioxide under alka-
line conditions have suggested a catalytic role for hydroxide ion
with a preferential binding to NO2, yielding ClO2ꢀ and NO3ꢀ as
products.34 If the assumption is valid, then in the presence of
[OHꢀ], the plot of k0/[ClO2] versus [OHꢀ] should give a
straight line. Such a linear curve should have an intercept equal to
k1 and a slope equal to kOHꢀ. Most likely, such a linear relation-
ship may not be observed at high concentrations of hydroxide,
when [OHꢀ] reaches stoichiometric proportions of the dye. The
plot of the second-order rate constant, k versus [OHꢀ] is illustrated
in Figure 4. Table 3 summarizes the calculated values of the second-
order rate constant. It is observed in Figure 4 that the y-intercept
value (k1) is very small, suggesting that in the absence of hydroxide
ion, the reaction is very slow or almost nil, which can be predicted
from the reported inert behavior of chlorine dioxide at acidic pH.35
The catalytic constant for the hydroxide catalyzed reaction in the pH
range of 6.0ꢀ7.5 was 4.0 ꢁ 109 Mꢀ2 sꢀ1 (Figure 4).
RAT LD50 190 mg kgꢀ1 37
.
Stoichiometric Equation. Using 1.5 ꢁ 10ꢀ3 M ClO2, the
stoichiometry of the reaction mixture was maintained with 1:1
and 1:5 ratios of amaranth and chlorine dioxide respectively. After a
30 min reaction, the residual amount reacted was determined, and
the amounts reacted were estimated. The stoichiometry was found
to be approximately 1:6 ((10%) of AMꢀ and ClO2. Thus, the
stoichiometric equation for the overall reaction can be written as
Oakes et al. during the oxidation of substituted arylazonaphthols
by hydrogen peroxide and Karkmaz et al. during the photocata-
lytic oxidation of amaranth have reported a quantitative conver-
sion of azo group to nitrogen.38,39 During the ozone initiated
oxidation of amaranth, no sulfate ion was detected as the product
immediately after the dye decolorization,40 but with extensive
mineralization of amaranth by ozone, sulfates were reported as
oxidation products.41 In the current study, immediately after the
decolorization of the dye with barium ion test, no sulfate was found
in the reaction mixture.
(c). Kinetic Salt Effect. From the studies of the reaction under
varied pH conditions, it is evident that the rate-limiting step in-
volves hydroxide ion. To confirm this assumption, the kinetic salt
effect on the reaction was investigated by measuring the reaction
rates at fixed concentrations of amaranth and chlorine dioxide,
but at varied ionic strengths. The plot of log k0 versus the square
root of ionic strength gave a positive slope (0.94), with a cor-
relation coefficient of R2 = 0.98, confirming that the rate-limiting
step involves like charges, possibly [OH
̅
] and AMꢀ. Assuming
that the main pathway of oxidation is OHꢀ catalyzed, and its re-
action order is unity, the overall third-order reaction coefficients
were calculated.
Reaction Scheme. The electrophilic attack by chlorine dioxide
on the nitrogen atom of AMꢀ abstracting an electron results
in the intermediate Int1. Int1 could resonate between three struc-
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The Journal of Physical Chemistry A
ARTICLE
tures. Subsequent attack by a hydroxyl ion on the carbon atom
bearing the nitrogen results in Int2 formation. Furthermore,
consecutive three electron abstractions from Int2 (two electrons
from nitrogen yielding N+ and one electron from the carbon
atom in between the two rings) results in the formation of the
dicationic Int3. The loss of a proton and rearrangement of the
electrons results in product P1 and a further intermediate, Int4.
The intermediate Int4 then undergoes two subsequent attacks
by hydroxyl ions resulting in Int5, followed by Int6 with con-
comitant loss of the N2 and NaSO3ꢀ ion. Then two sequential
abstractions of electrons by chlorine dioxide lead to the forma-
tion of naphthaquinone (P2).
Table 4. Elementary Reactions and Rate Coefficients Used
for Simulations
reaction
no.
elementary reaction
rate
AM ꢀ þ ClO2 f Int1 þ ClO2
ꢀ
(C1)
1.6 ꢁ 101 Mꢀ1 sꢀ1
1.68 ꢁ 108 Mꢀ1 sꢀ1
ꢀ
Int1 þ OH ꢀ f Int2
(C2)
(C3)
ꢀ
ꢀ
0.0 ꢁ 105 Mꢀ1
Int2 þ ClO2 f Int3a þ ClO2
5
sꢀ1
6.0 ꢁ 105
þ
-
M
ꢀ1 sꢀ1
Int3a þ ClO2 f Int3b þ ClO2
(C4)
(C5)
(C6)
(C7)
þ
2 þ
ꢀ
7.0 ꢁ 105 Mꢀ1 sꢀ1
4.32 ꢁ 105 Mꢀ1sꢀ1
Int3b þ ClO2 f Int3
þ ClO2
The structure of naphthaquinone was confirmed by gas chro-
1
matographyꢀmass spectrometry (GC-MS), H and 13C NMR
2 þ
þ
Int3
f Int4 þ P1 þ H þ
spectroscopy (Supporting Information Figures S1, S2, and S3).
The singlet peak at δ 6.9 are due to the olefinic proton, and the
resonances at δ 7.7 to δ 8.2 are due to the protons on the
aromatic ring. In the 13C NMR spectrum, the two carbonyl re-
sonances can be seen overlapping at δ 185 and the olefinic and
aromatic carbon resonances are seen between δ 125 and δ 140 .
The GC-MS shows a molecular ion peak at 158 for product P2 at
a retention time of 12.65 min. The loss of the carbonyl group
results in a fragmentation peak at m/z 130, and the loss of the
second carbonyl group can be seen at m/z 102 (Figure S2),
with the probable expansion of the six-membered ring to an
eight-membered ring. Analogous to the mechanism reported
by Oakes et al.,38 the identified oxidation product formation may be
explained by the mechanistic scheme illustrated in the following
scheme:
þ
Int4 þ OH ꢀ f Int5 þ N2
4 0.54 ꢁ 105 Mꢀ1
sꢀ1
þ
ꢀ
Int5 þ OH ꢀ f Int6 þ NaSO3
Int6 þ ClO2 f Int7 þ ClO2 þ H þ
Int7 þ ClO2 f P2 þ ClO2 þ H þ
(C8)
(C9)
(C10)
4.84 ꢁ 105 Mꢀ1 sꢀ1
5.17 ꢁ 106 Mꢀ1 sꢀ1
5.84 ꢁ 106 Mꢀ1 sꢀ1
ꢀ
ꢀ
be proposed as follows:
ClO2 þ AMꢀ f ClO2 þ Int1 ðvery slowÞ
ꢀ
2ꢀ
AMꢀ þ ClO2 þ OHꢀ f fHOAMClO2g
ðslow and rate limiting stepÞ
2ꢀ
fHOAMClO2g f ClO2 þ fHOAMgꢀ ðfastÞ
ꢀ
AMꢀ þ ClO2 f Int1 þ ClO2
ðR1Þ
ðR2Þ
ðR3Þ
ðR4Þ
ðR5Þ
ðR6Þ
ðR7Þ
ꢀ
ꢀ
Int1 þ OHꢀ f Int2
ꢀ
2þ
ꢀ
Int2 þ 3ClO2 f Int3 þ 3ClO2
Int3 f Int4 þ P1 þ Hþ
2þ
þ
Int4 þ OHꢀ f Int5 þ N2
þ
Int5 þ OHꢀ f Int6 þ NaSO3
ꢀ
Int6 þ 2ClO2 f P2 þ 2Hþ þ 2ClO2
ꢀ
Simulations. The computer simulation of the kinetic profiles
under varied reaction conditions were done with the elementary
reactions of the proposed mechanism (Table 4) and using Simkine
software.42,43 Estimated rate constants were adjusted using the
software until a best fit was obtained. The rate coefficients used
for the final simulations were optimized, and matching of the com-
puted curves with experimental profiles under varied [ClO2]0
conditions are shown in Figure 5. Table 4 summarizes the
experimentally determined values for eqs C1 and C2 used for
computation in the current study and the estimated rate coeffi-
cients (C3ꢀC10). The profiles of simulated curves matched well
with the experimental curves. The simulated and experimental
curves are illustrated in Figure 5. The simulated curves were
indicated by dashed lines and matched with the experimental curves.
Proposed Mechanism. Based on the reaction scheme pro-
vided and the stoichiometric results obtained, the mechanism can
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The Journal of Physical Chemistry A
ARTICLE
electrophilic abstraction, rather than by oxidative substitution or
addition. The rate-limiting step involved the formation of an
activated complex involving one each of ClO2, dye, and OHꢀ
ions. The catalytic constant for the hydroxyl ion was 4.0 ꢁ 109
Mꢀ2 sꢀ1 in the pH range 6.0ꢀ7.5. The energy parameters,
namely, energy and entropy of activations, were 50.0 kJ molꢀ1
and ꢀ658.7 J Kꢀ1 molꢀ1, respectively. The main oxidation
products were 1,2-dioxy-3-hyposulfite-8-sodium sulfite and 1,4-
naphthalenedione.
’ ASSOCIATED CONTENT
Supporting Information. 1H and 13C NMR data for
S
b
product P2, and mass spectroscopic data for the product P2. This mate-
rial is available free of charge via the Internet at http://pubs.acs.org.
Figure 5. Depletion of amaranth by chlorine dioxide. Experimental and
simulated curves for reaction of [AMꢀ]0 (7.0 ꢁ 10ꢀ5 M) with [ClO2]t
ꢁ 10ꢀ3/ M) (E1 = 0.025; E2 = 0.030,; E3 = 0.035) (experimental: solid
lines (curves E1ꢀE3); simulated: dashed lines (S1ꢀS3).
’ AUTHOR INFORMATION
Corresponding Author
*E-mail: jonnalagaddas@ukzn.ac.za.
’ ACKNOWLEDGMENT
The authors acknowledge the financial support received from
the University of KwaZulu-Natal and the National Research
Foundation, Pretoria, South Africa.
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