Oxidation of bromide by tert-butyl hypochlorite

Article abstract of DOI:10.1002/kin.20799

The kinetics of the oxidation reaction of bromide by tert-butyl hypochlorite (^tBuOCl) was studied at 25°C, ionic strength 0.5 M, and under isolation conditions. A stopped-flow spectrophotometer was employed for monitoring the reactions. Kinetic studies show that the reaction is first order with respect to [Br^-] and [^tBuOCl]. Linear dependences of the proton concentration, in perchloric acid medium, and the buffer solution concentration were found on the rate constant. The activation parameters were calculated using the Arrhenius and Eyring equations from the kinetic studies performed to analyze the influence of temperature on the rate constant. The results are consistent with a reaction mechanism of general acid catalysis. The catalytic constants were obtained for the oxidation of bromide by tert-butyl hypochlorite. The slope obtained for the Broensted relationship was 0.36.

Full text of DOI:10.1002/kin.20799

Oxidation of Bromide
by Tert-Butyl Hypochlorite
CRISTINA PASTORIZA, JUAN MANUEL ANTELO, JUAN CRUGEIRAS
Departamento de Qu´ımica F´ısica, Facultad de Qu´ımica, Universidad de Santiago de Compostela, 15782 La Corun˜a, Spain
Received 24 October 2012; revised 16 April 2013; accepted 20 April 2013
DOI 10.1002/kin.20799
Published online 26 August 2013 in Wiley Online Library (wileyonlinelibrary.com).
ABSTRACT: The kinetics of the oxidation reaction of bromide by tert-butyl hypochlorite (tBuOCl)
was studied at 25◦C, ionic strength 0.5 M, and under isolation conditions. A stopped-flow
spectrophotometer was employed for monitoring the reactions. Kinetic studies show that the
reaction is first order with respect to [Br−] and [tBuOCl]. Linear dependences of the proton
concentration, in perchloric acid medium, and the buffer solution concentration were found on
the rate constant. The activation parameters were calculated using the Arrhenius and Eyring
equations from the kinetic studies performed to analyze the influence of temperature on the
rate constant. The results are consistent with a reaction mechanism of general acid catalysis.
The catalytic constants were obtained for the oxidation of bromide by tert-butyl hypochlorite.
The slope obtained for the Bro¨nsted relationship was 0.36. ꢀC 2013 Wiley Periodicals, Inc. Int J
Chem Kinet 45: 629–637, 2013
This paper summarizes the kinetic results obtained
in the study of the reaction of Br− ions with tert-butyl
hypochlorite (tBuOCl) in acid medium. The reaction
was carried out using an excess of the Br− concen-
tration, at least 30 times higher than the BuOCl con-
centration. In these conditions, the reaction that takes
place is presented in Scheme 1.
To obtain kinetic information about the reaction pro-
cess and to discuss a possible reaction mechanism, a
set of experiments were designed. These experiments
allowed us to determine the reaction order with respect
to each one of the reagent concentrations and the influ-
ence of the concentrations of Br− and that of protons.
Also, the influence of ionic strength and temperature
was studied as well as the possible catalytic processes
when different buffer solutions were used to control
the proton concentrations.
INTRODUCTION
The chlorinating agents that have an atom of Cl+ are
important molecules because they have a high bacteri-
cidal capacity, so they are often used as agents for the
disinfection of drinking water [1–3]. However, their
use in water treatment can cause problems when other
chemical compounds are present, since these com-
pounds are usually highly reactive and have a high
capacity as oxidants [4–8]. The halides are an example
of compounds that can be easily oxidized by chlorinat-
ing agents, especially in acid medium [9–14]. Some
of these reactions involving halides can be used to as-
sess the amount of the chlorinating agent, but above all
they must be taken into account because the presence
of easily oxidizable species contributes to decrease the
Cl+ amount existing in the medium. This may be an
inconvenience during disinfection processes [1,15].
t
EXPERIMENTAL
Reagents
tBuOCl was obtained by passing a stream of chlo-
rine through a solution of NaOH, tert-butanol, and
Correspondence to: C. Pastoriza; e-mail: cinecris@gmail.com.
Supporting Information is available in the online issue at
www.wileyonlinelibrary.com.
C
ꢀ
2013 Wiley Periodicals, Inc.

630
PASTORIZA, ANTELO, AND CRUGEIRAS
−
sorption maximum of the Br3 species formed in the
reaction (λmax = 266 nm, εmax = 40,900 400 M−1
cm−1) [18]. This is the wavelength selected for moni-
toring of the reaction.
Scheme 1
water [16,17]. The product obtained was purified by
distillation and stored at 0◦C in a topaz flask to avoid
the decomposition by action of light. Its concentration
was determined iodometrically, and its purity was close
to 100%. The solution used in the kinetic experiments
Mixture of the Reaction
In previous studies, it has been found that this reaction
is very fast, so to perform the kinetic study it was nec-
essary to use a stopped-flow spectrophotometer. The
tBuOCl solution used in the kinetic experiments was
prepared every day, using acetonitrile as a solvent to
prevent decomposition and its hydrolysis. A system of
mixed asymmetric (1:25) was used with the purpose of
minimizing the solvent effect in the reaction medium.
The Br− concentration was always kept at least 30
times higher than the concentration of tBuOCl. All the
experiments were carried out at 25◦C, except a set of
experiments in which the influence of temperature was
studied. In all experiments, the ionic strength (I) was
maintained constant by adding the necessary volume
of 2 M NaClO4 solution.
The concentration of protons is a parameter that
has a great influence on the reaction rate, for this rea-
son it was necessary to maintain its concentration in
each experiment. This was achieved by adding a known
volume of a perchloric acid titrated solution or using
buffer solutions. In the latter case, it was also necessary
to study the possible catalytic effect of the species used
for the preparation of the buffer solutions.
t
was prepared every day, from BuOCl synthesized in
the laboratory, using acetonitrile as a solvent.
Aqueous solutions of NaBr, HClO4, and
NaClO4 were prepared using Aldrich commer-
cial products, A.C.S. reagent. Several buffer so-
lutions, prepared with Aldrich commercial prod-
ucts, A.C.S. reagent−, were used to control medium
pH: H3PO4/H2PO4
,
Cl2CHCOOH/Cl2CHCOO−,
ClCH2COOH/ClCH2COO−, CH3COOH/CH3COO−,
and CH3OCOOH/CH3OCOO−.
Instrumental
A spectrophotometer (Varian Cary 1 Bio) was used
for the previous studies. It was supplied with a ther-
mostated cell holder and was equipped with quartz
cells with 1-cm path length and a capacity of 3 cm3.
A stopped-flow spectrophotometer (Applied Photo-
physics DX.17MV) was used for monitoring of the
kinetics of the reactions. A system of mixed asymmet-
ric (1:25) was used with the purpose of minimizing
the solvent effect. The thermostate was used (Selecta
Frigiterm S-382) to ensure consistency in tempera-
ture of 0.1◦C. A workstation allowed obtaining the
absorbance–time data, which were collected by a com-
puter equipped with a program to analyze the data.
pH measurements were made with a pH meter PHM
82 Standard, Radiometer Copenhagen, equipped with a
combined electrode GK2401C, thermostated at 25◦C.
The pH meter was calibrated using standard solutions
supplied by Crison of pH 4.01, 7.00, and 9.26.
Reaction Order
Taking into account the stoichiometric equation for the
studied reaction, the rate equation can be written as
ꢀ
ꢁ
ꢀ
ꢁ
d BuOCl
d Br−3
t
v = −
=
t
t
d
d
ꢀ
ꢁ ꢀ ꢁ ꢀ
ꢁ
a
b
c
= k Br− H+ t BuOCl
(1)
As indicated above, all experiments were carried out
in conditions in which the concentrations of Br− and
H+ were practically constant during each experiment,
so that
RESULTS
Previous Studies: To Select the Wavelength
ꢀ
ꢁ
ꢁ
ꢀ
ꢁ
t
d BuOCl
d Br−3
v = −
=
To select the most suitable wavelength to monitor the
kinetics of the reaction, spectrophotometric studies of
all reagents and of the reaction mixture were carried
out in the range between 200 and 360 nm. The com-
parison of these studies has allowed us to confirm that
the absorption band of the reaction mixture shows a
maximum at 266 nm, which corresponds to the ab-
t
d
t
d
ꢀ
c
= kobs t BuOCl
(2)
where kobs represents the experimental rate constant.
To make sure that the reaction is first order
with respect to the BuOCl concentration, c = 1,
t
International Journal of Chemical Kinetics DOI 10.1002/kin.20799

OXIDATION OF BROMIDE BY TERT-BUTYL HYPOCHLORITE
631
Table I Rate Constants in Perchloric Acid Medium
T (◦C)
[tBuOCl] (M)
[H+] (M)
I (M)
[Br−] (M)
kobs (s−1
)
k2nd (× 103 M−1 s−1
)
25
25
25
25
25
25
25
25
25
25
25
25
25
25
25
15
25
35
45
5 × 10−5
5 × 10−5
5 × 10−5
5 × 10−5
5 × 10−5
5 × 10−5
5 × 10−5
5 × 10−5
5 × 10−5
5 × 10−5
5 × 10−5
5 × 10−5
5 × 10−5
5 × 10−5
5 × 10−5
5 × 10−5
5 × 10−5
5 × 10−5
5 × 10−5
0.01
0.01
0.01
0.01
0.0050
0.0075
0.010
0.015
0.025
0.050
0.02
0.02
0.02
0.02
0.02
0.01
0.01
0.01
0.01
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.0025
0.0050
0.010
26
50
102
254
38
55
50
90
139
1
4
4
7
2
4
1
4
3
10.4 0.4
10.0 0.8
10.2 0.4
10.2 0.3
7.6 0.4
11.0 0.8
10.0 0.2
18.0 0.8
27.8 0.6
0.025
0.0050
0.0050
0.0050
0.0050
0.0050
0.0050
0.0050
0.0050
0.0050
0.0050
0.0050
0.0050
0.0050
0.0050
0.0050
0.5
0.5
304 15
60.8
3
0.225
0.425
0.625
0.825
1.025
0.5
0.5
0.5
0.5
116
109
114
102
111
37
53
69
90
2
9
8
7
7
1
2
2
5
23.2 0.4
22
23
20
22
2
2
1
1
7.4 0.2
10.6 0.4
13.8 0.4
18
1
Table II Values of Activation Parameters for the
Reaction of Oxidation of tBuOCl by Br− in Perchloric
Acid Medium
the absorbance–time data were fitted to the inte-
grated equation written exponentially. In all cases, the
experiments were performed under isolation condi-
tions and the results were adjusted perfectly to this
equation:
Arrehnius Parameter
Eyring Parameter
A = (8 3) × 109 M–2 s–1 ꢀS= = (–63 3) J mol–1 K–1
Ea = (22.4 0.8) KJ mol–1 ꢀH= = (19.8 0.9) KJ mol–1
At = A∞ + (Ao − A ) · e−k
(3)
obst
∞
where A0, At, and A are the absorbance reading at
zero, t, and infinite times, respectively, and kobs is the
pseudo–first-order rate constant.
All the experiments were repeated between six and
10 times, and the deviation of the values obtained with
respect to the average value was less than 3%. The
mean values obtained for each experiment are pre-
sented in the tables.
∞
can be observed, and thus the reaction order with re-
spect to proton concentration is b = 1.
On the other hand, we have studied the influence
of the ionic strength (I) on the rate constant in acid
medium. The results indicate that the rate constant is
independent of the ionic strength under the conditions
studied (Table I). This seems to indicate that at least
one nonionic species must be involved in the transition
state.
A study of the influence of temperature on the rate
constant was performed to complete the study of the re-
action of Br− with tBuOCl in perchloric acid medium.
This study was carried out at 15, 25, 35, and 45◦C.
The results listed in Table I show an increase in the
rate constant with increasing the temperature. These
experimental results are in compliance with the Arrhe-
nius equation and the equation derived from the Eyring
theory activated complex. This allows us to obtain the
activation parameters for the reaction in perchloric acid
medium (Table II).
Reaction in Perchloric Acid
The results summarized in Table I show an increase in
the rate constant by increasing the Br− concentration,
so that the reaction order with respect to the bromide
concentration is a = 1. This table also includes the
second-order rate constant, k2nd = kobs/[Br−].
Another group of experiments allows us to confirm
the increase in the rate constants kobs and k2nd with
increasing the proton concentration (Table I). A linear
relationship between kobs and the proton concentration
International Journal of Chemical Kinetics DOI 10.1002/kin.20799

632
PASTORIZA, ANTELO, AND CRUGEIRAS
Table III Influence of the Concentrations of Br− and tBuOH, in the Presence of Buffer Solution
Buffer Solution
[tBuOH] (M) pH [BR–] (M) kobs (s–1
)
k2nd (× 103 M–1 s–1
)
[H3PO4/H2PO4 –]T = 0.05 M, I = 0.425 M
–
–
1.85
1.86
1.85
2.58
2.57
2.43
2.40
2.09
2.13
2.14
2.12
2.11
2.12
2.12
2.16
0.005
0.015
0.030
0.005
0.010
0.025
0.035
0.005
0.005
0.005
0.005
0.005
0.005
0.005
0.005
71
224
2
4
14.2 0.4
14.9 0.3
13.5 0.7
4.6 0.2
4.7 0.1
4.6 0.2
4.60 0.06
7.6 0.4
9.0 0.2
8.2 0.2
8.2 0.2
7.8 0.4
9.6 0.4
8.9 0.2
7.8 0.4
–
404 20
[ClCH2COOH/ClCH2COO–]T = 0.2 M, I = 0.4 M
[H3PO4/H2PO4–]T = 0.2 M, I = 0.4 M
–
–
–
–
23
47
116
161
38
45
41
41
39
1
1
4
2
2
1
1
1
2
2
0.010
0.025
0.050
0.075
0.10
0.25
0.35
0.50
48
44.3 0.8
39
2
T = 25◦C, [tBuOCl] = 5 × 10−5 M.
The calculated enthalpy value shows that the pro-
cess corresponds to a chemical control for the reaction.
The negative activation entropy value indicates that the
transition state is more orderly than the reactants.
Kinetics Studies in the Presence of Buffer
Solutions
To study the influence of pH on the kinetics of the
oxidation reaction of bromide by tert-butyl hypochlo-
rite, the kinetic studies were carried out using different
buffer solutions to control the proton concentration in
the medium.
The results obtained with different concentrations
of bromide, for each buffer solution, are included in
Table III. A linear increase in the rate constant with an
increase in the bromide concentration can be observed.
This confirms that the reaction is first order with respect
to the bromide concentration.
Figure 1 Influence of [ClCH2COOH/ClCH2COO−]T on
kobs. T = 25◦C, I = 0.5 M, [tBuOCl] = 5 × 10−5 M, [Br−] =
5 × 10−3 M. •: pHm = 1.85, ◦: pHm = 2.31 (dash line), ꢀ:
pHm = 2.68, ꢁ: pHm = 3.27, ꢂ: pHm = 3.65.
Another group of experiments allows us to confirm
t
that the presence of the BuOH does not significantly
alter the rate constant (Table III).
In each experiment, we can observe that the rate
constant increases with increasing the buffer solution
concentration. This suggests the possible existence
of a catalysis process. This catalytic effect was ob-
served in all buffer solutions used, and the linear rela-
tionships obtained in each experiment between kobs
and the buffer solution concentration are shown in
Figs. 1–4.
Table IV summarizes the results obtained by least-
squares analysis of k2nd versus [HAi]T for all se-
ries of experiments carried out with different buffer
solutions.
Influence of the Buffer Solution
Concentration
A set of experiments were carried out to study
the effect of different concentrations of buffer so-
lutions on the reaction medium. In these experi-
ences, the pH of the medium was modified using dif-
ferent buffer solutions: Cl2CHCOOH/Cl2CHCOO−,
ClCH2COOH/ClCH2COO−, CH3COOH/CH3COO−,
and CH3OCOOH/CH3OCOO−.
International Journal of Chemical Kinetics DOI 10.1002/kin.20799

OXIDATION OF BROMIDE BY TERT-BUTYL HYPOCHLORITE
633
Figure 4 Influence of [Cl2CHCOOH/Cl2CHCOO−]T on
kobs. T = 25◦C, I = 0.5 M, [tBuOCl] = 5 × 10−5 M, [Br−] =
5 × 10−3 M. •: pHm = 1.71.
Figure 2 Influence of [CH3OCOOH/CH3OCOO−]T on
kobs. T = 25◦C, I = 0.5 M, [tBuOCl] = 5 × 10−5 M, [Br−] =
5 × 10−3 M. •: pHm = 3.59.
MECHANISM AND DISCUSSION
t
When the solutions of BuOCl and Br− are mixed,
−
the formation of the Br3 species takes place. Experi-
mentally, the reaction was found to be first order with
respect to the concentration of each one of the reagents.
Moreover, the possibility that the reaction may be in
equilibrium was discarded because the rate constant is
not affected by the change in the tBuOH concentration.
The reaction that occurs is a redox process where
the Cl+, of the tBuOCl, will be reduced to Cl− while
the bromide ion (Br−) is oxidized to Br0.
The literature indicates that, for other similar reac-
tions, the formation of bromine may take place through
the formation of the BrCl intermediate compound,
which in excess of Br− quickly leads to the forma-
tion of Br2 [11,18,19]. On the other hand, according
Figure 3 Influence of [CH3COOH/CH3COO−]T on kobs
.
T = 25◦C, I = 0.5 M, [tBuOCl] = 5 × 10−5 M, [Br−] = 5
× 10−3 M. •: pHm = 3.94, ꢂ: pHm = 4.62, ꢀ: pHm = 5.37.
Table IV Catalytic Constants for the Reaction of Oxidation of Br− by tBuOCl
Buffer
pHm fHA Intercept (× 103 M–1 s–1) Slope (× 103 M–1 s–1
)
r
kHA (× 103 M−2 s−1
)
i
i
Cl2CHCOOH/Cl2CHCOO– 1.71 0.26
pKa = 1.26
17.2 0.6
4.4 0.6
3.8 0.2
2.2 0.2
0.6 0.1
0.22 0.02
47.0 2.0
26
26.6 0.6
13
5.6 0.2
2.8 0.1
0.9952
0.9819
0.9988
0.9878
0.9941
0.9976
181
31
8
3
1.85 0.85
2.31 0.66
2.68 0.45
3.27 0.18
3
40.3 0.9
1
29
31
35
2
1
1
ClCH2COOH/ClCH2COO– 3.65 0.08
pKa = 2.60
CH3OCOOH/CH3OCOO–
3.59 0.46
3.94 0.81
4.62 0.48
5.39 0.13
0.24 0.06
0.12 0.01
0.06 0.08
0.024 0.004
4.2 0.2
4.54 0.04
3.0 0.2
0.9936
0.9997
0.9828
0.9988
9.1 0.4
5.60 0.05
pKa = 3.53
6
2
CH3COOH/CH3COO–
0.63 0.02
4.8 0.1
pKa = 4.58
T = 25◦C, I = 0.5 M.
International Journal of Chemical Kinetics DOI 10.1002/kin.20799

634
PASTORIZA, ANTELO, AND CRUGEIRAS
This includes three terms: (1) a reaction in aqueous
solution without a catalyst (kH O[H2O] =k0), (2) catal-
2
ysis due to the presence of protons (kH O+), and (3)
3
catalysis due to the acid specie of the buffer solution
(kHA ).
i
Given that the study has been performed under isola-
tion conditions, tBuOCl being the limiting reactant and
making a material balance, we can express the rate con-
stants kobs and k2nd as a function of the buffer concen-
tration total ([HAi]T) and of the acid fraction undisso-
ciated from the buffer solution (fHAi = [HAi]/[HAi]T):
Scheme 2
ꢀ
ꢁ
ꢀ
ꢁ ꢀ
ꢁ
kobs = k0 Br− + kH O H3O+ Br−
+
3
ꢀ
ꢁ
to the literature, when Br2 is present in solution along
Br−, the Br3− species will be formed in a fast equilib-
rium process that is dissociated from the formation of
Br3− (Ke = 16.1 at 25◦C, I = 1.0 M) [18].
+ kHA fHA [HAi]T Br−
(6)
i
i
ꢀ
ꢁ
k2nd = k0 + kH O H3O+
+
3
+ kHA fHA [HAi]T
(7)
The experimental results show an increase in the rate
constant with increasing the concentration of buffer so-
lution, when the concentrations of the hydrogen ion and
bromide ion remain constant. Also, an increase in the
rate constant with increasing the proton concentration
is observed when the pH of the medium is modified.
This behavior is consistent with that observed in other
reactions in which the process of general acid–base
catalysis takes place [11].
i
i
These expressions are consistent with the experimental
results since they allow us to justify: (1) the first order
with respect to the tBuOCl concentration, (2) the first
order with respect to the Br− concentration, (3) the
linear dependence between the rate constant and the
proton concentration, and (4) the increase in the rate
constant with increasing the buffer concentration.
According to the above, the proposed mechanism
for this oxidation process of the Br− ion is shown in
Scheme 2.
Calculation of Catalytic Constant
Reaction in Perchloric Acid Medium. When the pro-
ton concentration of the medium is controlled using a
perchloric acid solution, the general equation for the
rate experimental constant (Eq. (7)) is simplified as
The estimated value for the dissociation constant of
the tBuOHCl+ is approximately 102 [20], so that under
the experimental conditions, [H+] < 0.05 M, more than
99.99% of the tert-butyl hypochlorite is present as neu-
ꢀ
ꢁ
+
t
tral species (tBuOCl). Therefore, BuOCl will be the
follows: k2nd = k0 + k
H3O+
.
H3O
The results summarized in Table I allow us to prove
that a linear relationship exists between k2nd and the
proton concentration, as shown in Fig. 5.
species involved in the slow step of the reaction. From
this mechanism, we obtained a general equation for the
rate reaction that is a sum of terms, in which the differ-
ent species that can act as catalysts of the process are
included:
The value of the intercept allows us to calculate the
term: kH O[H2O] = k0, k0 = (0.0 0.1) × 104 M−1
2
s
−1. This term in a strongly acidic medium is almost
ꢀ
ꢁ
negligible compared to the term kH O+, but it may be
3
t
d BuOCl
important if the pH increases. However, the studies
carried out using buffer solutions will make it possible
to verify the existence of a process of acid catalysis,
so the reaction rate is to be strongly influenced by the
catalytic constant.
v = −
t
d
ꢂ
ꢀ
ꢁ ꢀ
ꢁ
=
kHA [HAi] Br− t BuOCl
(4)
i
i
The slope value corresponds to the proton catalytic
The presented expression of this rate equation can
be written as
constant, kH O+, and its value can be calculated from
3
the experimental results:
ꢀ
ꢁ ꢀ
ꢁ ꢀ
ꢁ
v = kH O [H2O] Br− t BuOCl
kH O = (1.19 0.05) × 106 M−2 s−1
+
2
3
ꢀ
ꢁ ꢀ
ꢁ ꢀ
ꢁ
+ kH O H3O+ Br− t BuOCl
+
3
Figure 6 shows the values of the second-order rate con-
stant, k2nd against pH. The values presented in Table I
ꢀ
ꢁ
+ kHA [HAi] Br− t BuOCl
(5)
i
International Journal of Chemical Kinetics DOI 10.1002/kin.20799

OXIDATION OF BROMIDE BY TERT-BUTYL HYPOCHLORITE
635
Figure 7 Plot of k2nd against [H+] in the oxidation of Br−
Figure 5 Influence of [H+] on k2nd (Table I).
by tBuOCl.
acids present in the reaction medium. The data shown
in Figs. 1–4 allow us to prove for each buffer solu-
tion the existence of a linear relationship between k2nd
and the total buffer concentration. Table IV presents
the results obtained for the intercepts and slopes by
adjustments of k2nd versus [Buffer]T.
The slopes of these linear relationships (slope =
kHA fHA ) allow us to calculate the catalytic con-
i
i
stant of each of the acids used for the preparation
of buffer solutions. These values are also included
in Table IV. In this table, we can observe an in-
crease in the catalytic constant kHAi with increasing the
acid strength in the following sequence: Cl2CHCOOH/
Cl2CHCOO−
>
ClCH2COOH/ClCH2COO−
>
CH3OCOOH/CH3OCOO− > CH3COOH/CH3COO−.
These results are in concordance with those obtained
in previous studies [21,22].
Figure 6 Plot of k2nd against pH in the oxidation of Br−
by tBuOCl (Table I).
are also included in this figure, which shows that the
obtained profile is exponential. This suggests that the
second-order rate constant tends to be constant at high
pH values. This value would correspond to the value of
The values of the catalytic constants are summa-
rized in Table V. The results presented in this ta-
ble, for the buffer solutions of dichloroacetic acid and
methoxyacetic acid, were obtained by varying the con-
centration of buffer at only one pH value. In the case
of the buffer solutions of monochloroacetic acid and
acetic acid, the values included in this table are the av-
erage values of the rate constants obtained in five and
three series of different experiments, respectively.
Moreover, the intercept values summarized in
Table IV correspond to the rate constant at the zero
buffer concentration. This can be interpreted as the
rate constant in the absence of buffer at that pH value,
which may be analyzed together with the rate constants
obtained in perchloric acid medium.
the term kH O[H2O]. As will be seen later, the value of
2
k2nd approaches zero when considering the results ob-
tained with buffer solutions, which allows us to extend
the pH range studied.
Reaction in the Presence of Buffer Solutions. A sec-
ond set of experiments was carried out using differ-
ent buffer solutions in which the weak acids used
to obtain these solutions were H3PO4, Cl2CHCOOH,
ClCH2COOH, CH3COOH, and CH3OCOOH. Thus,
the reaction was studied in the range of pH between
1.26 and 4.58.
The experimental results indicate that the reaction
rate depends on the nature of the species used for the
preparation of the buffer solutions and its concentra-
tion. Also, the reaction rate depends on the pKa of the
Figure 7 shows the values of the rate constant k2nd
obtained in the experiments carried out in perchloric
acid medium along with the intercepts obtained in
the study of the influence of the buffer concentration
(Table IV). The good linear relationship is shown
International Journal of Chemical Kinetics DOI 10.1002/kin.20799

636
PASTORIZA, ANTELO, AND CRUGEIRAS
Table V Catalytic Constants for the Reaction of Oxidation of Br− by tBuOCl
a
Species
pKa
kHA ( × M−2 s−1
)
kHOCl (× M–2 s–1
)
i
H3O+
–1.74
1.26
2.60
3.53
4.50
15.52
(1.15 0.05) × 106
(1.81 0.08) × 105
(3.3 0.5) × 104
(9.1 0.4) × 103
(5.8 0.6) × 103
–
(1.32 0.03) × 106
Cl2CHCOOH/Cl2CHCOO–
ClCH2COOH/ClCH2COO–
CH3OCOOH/CH3OCOO–
CH3COOH/CH3COO–
H2O
–
(6.11 0.34) × 104
–
(2.09 0.01) × 104
27.9 5.5
aData obtained from [11] for the reaction Br– + HOCl.
in this figure. The slope value obtained from the
least-squares adjustments allows to obtain a value
+
for the catalytic constant, kH O = (1.15 0.05) ×
3
106 M−2 −1, which is completely in accordance with
s
that obtained in the experiments performed in perchlo-
ric acid medium.
Bro¨nsted Relationship
The acid dissociation constant (Ka) and the rate con-
stants for the reactions catalyzed by a buffer solution
(kHA) can be correlated to the Bro¨nsted equation [23],
which is shown in the following equation:
ꢃ
ꢄ
α
kHA
p
Ka q
p
= GA
(8)
Figure 8 Bro¨nsted plots for the reactions of oxidation of
t
Br− by BuOCl or by HOCl. •: Data obtained in this work
t
(Br− + BuOCl). ◦: Data from [11] (Br− + HOCl).
where p is the number of acid equivalent positions in
the conjugate acid, q is the number of basic equiva-
lent positions in the conjugate base, GA is the Bro¨nsted
proportionality constant, and α is the measurement of
the proton transfer in the transition state. The α values
are often in the range of 0 < α < 1 for different reac-
tions in aqueous solution. This equation corresponds
to a free-energy relationship between the free ener-
gies of activation for proton transfer reactions and free
energies of acid ionization.
state is closer to the reactants or products. The value
found reflects that the proton transfer is 36%.
It is important to consider that in these processes
t
the proton HA cannot be transferred to BuOCl in a
preequilibrium step, because otherwise we would have
a specific catalytic process, where the rate constant de-
pends only on the concentration of protons and not on
the buffer concentration. This means that the transition
state must contain three species: HA, tBuOCl, and Br−.
The compliance of this relationship can be verified
if the Bro¨nsted equation is expressed in the logarithmic
form.
t
The oxygen atom of the species BuOCl accepts a
The values of α and the proportionality constant of
Bro¨nsted (GA) can be obtained from the values of the
slope and the intercept obtained by least squares in the
log (kHA/p) versus (log (Kaq/p)), respectively. Figure 8
shows that the results obtained in this study and the
values for the reaction of Br− with HOCl match with
the literature [11].
proton of HA, whereas the Cl+ is transferred to Br−.
To reach this transition state, it is not necessary that a
trimolecular collision takes place; it is only needed that
the species HA approaches to the weakly associated
species (tBuOCl+Br−) or, perhaps, an approach of the
Br− ion to species (AHtBuOCl) occurs, always taking
into account that the proton transfer cannot be fully
carried out before the association of the Cl+ and Br−
(see Scheme 3).
The values for the intercept and slope obtained in
this representation are 5.3
0.1 and 0.36
0.05,
respectively. The term α, corresponding to the slope
of the Bro¨nsted, is often interpreted as a measure of
the extent of proton transfer in the transition state. The
reasoning is the interpretation of whether the transition
In this transition state, there are five pairs of elec-
trons around the chlorine atom. This situation is not
unusual, taking into account that it appears in a sta-
ble species as Cl3−. The valence shell of electron
International Journal of Chemical Kinetics DOI 10.1002/kin.20799

OXIDATION OF BROMIDE BY TERT-BUTYL HYPOCHLORITE
637
BIBLIOGRAPHY
1. White, G. C. The Handbook of Chlorination, 3rd ed.;
Van Nostrand Reinhold; New York, 1992; 889–934.
2. Morris, J. C. Water Chlorination Environmental Impact
and Health Effects; Ann Arbor Science: Ann Arbor, MI,
1978; Vol 2, pp. 879–890.
3. Jolley, R. L.; Carpenter, J. H. Water Chlorination Envi-
ronmental Impact and Health Effects; Ann Arbor Sci-
ence: Ann Arbor: MI, 1983; Vol. 4, Chap. 1, pp. 3–47.
4. Jensen, J. S.; Lam, Yiu-Fai; Helz, G. R. Environ Sci
Technol 1999, 33, 3568–3573.
Scheme 3
pair repulsion theory predicts a bond between O, Cl+,
and Br−.
5. Yang, X.; Fan, Chihhao; Shang, C.; Zhao, Q. Water Res
2010, 22, 2691–2702.
6. Donnermair, M. M.; Blantchley, E. R. Water Res 2003,
37, 1557–1570.
7. Liu, S.; Sung, G. Ind Eng Chem Res 2009, 48, 613–618.
8. Rees, M. D.; Pattison, D. I.; Davies, M. J. Biochem J
2005, 391, 125–134.
9. Deborde, M.; von Gunten, U. Water Res 2008, 42, 13–
51.
10. Kumar, K.; Day, R. A.; Margerum, D. W. Inorg Chem
1986, 25, 4344–4350.
11. Kumar, K.; Margerum, D. W. Inorg Chem 1987, 26,
2706–2711.
12. Johnson, D. W.; Margerum, D. W. Inorg Chem 1991,
30, 4845–4851.
13. Nagy, J. C.; Kumar, K.; Margerum, D. W. Inorg Chem
1988, 27, 2773–2780.
CONCLUSIONS
There is information in the literature about the oxida-
tion reaction of Br− with HOCl, but there is no in-
formation for the kinetics of the oxidation reaction of
Br− with tBuOCl. The results obtained in the present
study of the oxidation reaction of Br− with BuOCl
t
have allowed us to prove that its kinetic behavior is
totally identical to the behavior found in the reaction
of oxidation of Br− with HOCl.
In general, the results can be summarized as
follows:
– first order with respect to the concentration of the
Br− ion,
– first order with respect to the concentration of
14. Bichsel, Y.; Von Gunten, U. Environ Sci Technol 2000,
34, 2784–2791.
tBuOCl,
– the rate constant is not affected by the variation
of the concentration of tBuOH, so the possibility
that the reaction is in equilibrium is discarded,
– complex dependence of the rate constant with
respect to the H+ concentration, and
– the reaction catalyzed by acids and bases, so that
the species used as pH regulators will act as cat-
alysts for the reaction. The catalytic constants
obtained are adjusted to the Bro¨nsted equation.
15. Margerum, D. W.; Huff Hartz K. E. J Environ Monit
2002, 4, 20–26.
16. Teeter, H. M.; Bell, E. W. Org Synth 1963, 4, 125.
17. Walling, C.; Jacknow, B. B. J Am Chem Soc 1960, 82,
6108–6112.
18. Wang, T. X.; Kelley M. D.; Cooper, J. N.; Beckwith, R.
C.; Margerum, D. W. Inorg Chem 1994, 33, 5872–5878.
19. Anbar, M.; Ginsberg, D. Chem Rev 1954, 54, 925–958.
20. Anbar, M.; Guttmann, S.; Rein, R. J Am Chem Soc
1959, 81, 1816–1821.
21. Antelo, J. M.; Arce, F.; Campos, J.; Parajo´, M. Int J
Chem Kinet 1996, 28, 391–396.
22. Antelo, J. M.; Arce, F.; Crugeiras, J.; Miraz, C.; Parajo´,
M. Gazz Chim Ital 1997, 127, 355–360.
23. Maskill, H. The Physical Basis of Organic Chemistry;
Oxford University Press: Oxford, UK, 1985; Chap. 10.
SUPPORTING INFORMATION
Tables presenting data corresponding to the buffer so-
lution concentration, for each one of the buffers used in
this work, are included in the Supporting Information.
International Journal of Chemical Kinetics DOI 10.1002/kin.20799