Kinetic evidence for accumulation of stoichiometrically significant amounts of H2I2O3 during the reaction of I- with IO3-

Article abstract of DOI:10.1021/jp000271w

The kinetics of the Dushman reaction, 5I^- + IO₃^- + 6H⁺ → 3I₂ + 3H₂O, is considerably more complex in aqueous, acidic media than is the behavior of the analogous reaction of Br^- with BrO₃^-. The latter reaction is first order in both [Br^-] and [BrO₃^-] and is believed to proceed via a series of simple oxygen-atom transfers, there being no evidence for the intermediacy of polybromine species such as H₂Br₂O₃. In contrast, the [I^-] order in the Dushman reaction is reported variously to be one or two, depending upon experimental conditions, especially [IO₃^-]₀/[I^-]₀, and the kinetics of this reaction is generally very complex with the intermediacy of H₂I₂O₃ having been suggested by several authors. Stopped-flow kinetic measurements of the Dushman reaction are reported here with excess [KIO₃]₀ ≈ 40[KI]₀/5 and [HClO₄]₀ ≥ 140 × 6[KI]₀/5 in which the observed [I^-] order is cleanly two, but saturation kinetics is observed for the pseudo-second-order rate constant as [KIO₃]₀ and [HClO₄]₀ increase. This observation is rationalized by a mechanism assuming the intermediacy of H₂I₂O₃ at concentrations significant compared to those of I^-, requiring that the kinetic equations be constrained by the iodide-ion mass balance. It is suggested that H₂I₂O₃ is a kinetically and stoichiometrically significant species in the Dushman reaction, while H₂Br₂O₃ is not in the analogous bromine system, because of the greater atomic radius, polarizability, and ability to expand its valence shell of iodine as compared to bromine.

Full text of DOI:10.1021/jp000271w

J. Phys. Chem. A 2000, 104, 5269-5274
5269
Kinetic Evidence for Accumulation of Stoichiometrically Significant Amounts of H2I2O3
-
during the Reaction of I- with IO3
Jesu´s Alberto Agreda B.*
Departamento de Quimica, UniVersidad Nacional de Colombia, Santafe´ de Bogota´, Colombia
Richard J. Field* and Nancy J. Lyons
Department of Chemistry, The UniVersity of Montana, Missoula, Montana 59812
ReceiVed: January 21, 2000; In Final Form: March 2, 2000
-
The kinetics of the Dushman reaction, 5I- + IO3 + 6H+ f 3I2 + 3H2O, is considerably more complex in
-
aqueous, acidic media than is the behavior of the analogous reaction of Br- with BrO3 . The latter reaction
-
is first order in both [Br-] and [BrO3 ] and is believed to proceed via a series of simple oxygen-atom transfers,
there being no evidence for the intermediacy of polybromine species such as H2Br2O3. In contrast, the [I-]
order in the Dushman reaction is reported variously to be one or two, depending upon experimental conditions,
-
especially [IO3 ]0/[I-]0, and the kinetics of this reaction is generally very complex with the intermediacy of
H2I2O3 having been suggested by several authors. Stopped-flow kinetic measurements of the Dushman reaction
are reported here with excess [KIO3]0 ≈ 40[KI]0/5 and [HClO4]0 g 140 × 6[KI]0/5 in which the observed
[I-] order is cleanly two, but saturation kinetics is observed for the pseudo-second-order rate constant as
[KIO3]0 and [HClO4]0 increase. This observation is rationalized by a mechanism assuming the intermediacy
of H2I2O3 at concentrations significant compared to those of I-, requiring that the kinetic equations be
constrained by the iodide-ion mass balance. It is suggested that H2I2O3 is a kinetically and stoichiometrically
significant species in the Dushman reaction, while H2Br2O3 is not in the analogous bromine system, because
of the greater atomic radius, polarizability, and ability to expand its valence shell of iodine as compared to
bromine.
Introduction
The kinetics of the reaction of IO3- with I- in aqueous, acidic
media
rate expressions containing up to three terms in unbuffered
media. All three terms are supposed by these authors to be first
order in [IO3-] and second order in [H+]. However, the kinetic
dependence on [I-] is very complex, with two of these terms
being first order and second order, respectively, in [I-], and
IO3- + 5I- + 6H+ a 3I2 + 3H2O
(1)
-
the third term being first order in both [I-] and [I3-]. The [I3
]
dependence arises from reaction 3, which is not a significant
was first investigated by Dushman1 in 1904 and proceeds far
to the right in strongly acid media. The most complete
experimental investigation of reaction 1 seems to be that of Abel
and Stadler2 in 1926 followed by that of Abel and Hilferding3
in 1928. Interest in the oscillatory Bray-Liebhafsky reaction4-7
led to a resurgence of work8-11 on this system in the 1970s,
including a review and simulation of existing data by Liebhafsky
and Roe11 in 1979. Such activity continues within the context
of oscillating chemical reactions.12-14 Some recent work has
been motivated by a need to understand the aqueous chemistry
of radioactive iodine species in the event of a nuclear reactor
containment breach.15
-
I2 + I- a I3
(3)
perturbation at the very low [I-] present here, for which [I3-]/
[I2] seldom exceeds 0.05 (assuming17 K3 ) 740 M-1), and then
only early in the reaction when [I2] is relatively low and [I-] is
relatively high. The kinetics of reaction 1 is even more complex
in buffered media, leading to yet a fourth term in the rate
expression.8,9,11
The kinetic complexity of the Dushman reaction was ascribed
by Bray4 in 1930 and later by Liebhafsky and co-workers18-20
to the intermediacy of H2I2O3 or its anhydride, I2O2. This species
has been invoked by later workers10,14,20,21 in more complex
The equivalent reaction with bromine seems kinetically clean
and simple.16
-
reactions, as have the related intermediate species IO3H2SO3
suggested by Rabai et al.12 in the IO3--SO32--Fe(CN)6
4-
BrO3- + 5Br- + 6H+ a Br2 + 3H2O
(2)
system, and I2O suggested by Urbansky et al.15 in regard to the
disproportionation of HOI, as well as by Schmitz14 and by Kolar-
Anic22-24 and co-workers in regard to the mechanism of the
Bray-Liebhafsky reaction.4-7 Schmitz14c very recently also
carried out a systematic analysis of then existing kinetic results
on reaction 1. We report here stopped-flow kinetic evidence
for both the participation and the accumulation of stoichiomet-
Reaction 2 is first order in both [Br-] and [BrO3-] and appears
to occur via a series of rapid oxygen-atom transfers without
the intermediacy of polybromine species such as H2Br2O3.
However, the kinetics of the Dushman reaction is complex.
Liebhafsky and Roe11 attempted to describe the kinetics with
10.1021/jp000271w CCC: $19.00 © 2000 American Chemical Society
Published on Web 05/13/2000

5270 J. Phys. Chem. A, Vol. 104, No. 22, 2000
Agreda B. et al.
rically significant (compared to [I-]) amounts of H2I2O3 in the
Dushman reaction when [KIO3]0 . [KI]0/5.
Experimental Section
All reagents were made up using deionized water. Experi-
ments were carried out in 0.0968-0.489 M HClO4 (Fisher 69-
72% Metal Grade) aqueous media. Stock solutions were made
up using EM Science GR grade KI without further purification.
Baker Analyzed Reagent grade KIO3 was slurried in methanol
to remove I- and then used to prepare stock solutions. Kinetic
experiments were carried out spectrophotometrically, usually
at 460 nm near the absorption maximum of I2, at room
temperature, 24 ( 0.5 °C, using an Aminco-Morrow stopped-
flow apparatus upgraded by On-Line Instrument Systems, Route
2, Jefferson, GA 30549. Digital data-acquisition methods were
used and results were manipulated using standard PC methods
and software. The stoichiometric iodide remaining at time t,
[I-]s, was calculated from the measured instantaneous [I2]t
5
absorbance using the relation [I-]s ) /3([I2]∞ - [I2]t). In the
mechanism described below, stoichiometric iodide, [I-]s may
be distributed between I-, [I-]t, and H2I2O3, [H2I2O3]t, i.e., [I-]s
) [I-]t + [H2I2O3]t.
Results
All experiments were carried out with [KIO3]0 ≈ 40 [KI]0/5,
[HClO4]0 ≈ [H+]0 g 140 × 6[KI]0/5, and [HClO4]0
>
20[KIO3]0, ensuring isolation kinetics in [I-]s at constant
[HClO4]0, and [HIO3]0, the last concentration being calculated
from [HIO3]0 ) (Ka + [KIO3]0 + [HClO4]0 - ((Ka + [KIO3]0
+ [HClO4]0)2 - 4[HClO4]0[KIO3]0)1/2)/2. The value Ka0 ) [IO3-]-
[H+]/[HIO3] ) 0.157 at 25 °C and infinite dilution (µ f 0) is
given by Strong and Pethybridge.25 It is expected26 that the value
of γ( will be nearly constant and about 0.7 for a one-to-one
electrolyte such as HIO3 in the range 0.1-0.5 M. Thus, we
0
2
calculate Ka ) Ka /γ( ) 0.32. The kinetics under these
conditions and in the concentration ranges 0.000 058 M e [KI]0
e 0.000 576 M, 0.000 448 M e [KIO3]0 e 0.004 48 M, and
0.0968 M e [HClO4]0 e 0.489 M is found to be second order
(after a very short induction period in some of the most rapid
experiments; e.g., see Figure 6) in [I-], with no first-order
component, i.e., eq 4. Figure 1 shows typical absorbance data
Figure 1. (Lower) typical stopped-flow kinetic data (500 points)
showing the average absorbance × 102 due to I2 at 460 nm in 10
duplicate experiments vs time (s). (Upper) second-order plot of (1/
[I-]s)/105 (M-1) vs time (s). The stoichiometric [I-]s is calculated from
d[I-]
1
5
-
s ) kexp[I-]s2
(4)
dt
5
the instantaneous [I2]t according to [I-]s ) /3([I2]∞ - [I2]t). Slope )
as well as the corresponding second-order plot of 1/[I-]s, which
is linear to nearly 99% completion. The values of kexp reported
here are the slopes of plots such as that displayed in Figure 1
divided by 5 in order to take into account the stoichiometric
factor in eq 4.
(4.6 ( 0.5) × 104 M-1 s-1 and kexp ) slope/5 ) (9.2 ( 1) × 103 M-1
s-1. [HClO4]0 ) 0.289 M, [KIO3]0 ) 0.000 448 M, [HIO3]0 ) 0.000 212
M, [KI]0 ) 0.000 058 M.
Figures 2 and 3 and the mechanism described below suggest
Pseudo-second-order rate constants, kexp, were determined
with varying [KIO3]0 and [HClO4]0 and display saturation
kinetics as these concentrations increase. Figure 2 shows five
series of kexp plotted vs ([HIO3]0[HClO4]0) with each series
having constant [HClO4]0 but varying [KIO3]0. The series at
lowest [HClO4]0 is nearly linear with no unequivocal evidence
of an intercept, indicating a single-term form of kexp. However,
saturation (negative) curvature appears at higher ([HClO4]0
[HIO3]0) in the series at higher values of [HClO4]0. Figure 3
shows five series of kexp plotted vs ([HIO3]0[H+]0) with each
series having constant [KIO3]0 but varying [HClO4]0. The series
at lowest [HIO3]0 again is nearly linear without unequivocal
evidence of an intercept, but the series at higher [HIO3]0 show
substantial saturation deviation as [HClO4]0 increases.
eqs 5-7.
d[I-]s
k[HIO3]0[HClO4]0
(1 + k*[HIO3]0[HClO4]0)2
1
5
-
)
[I-]s2 ) kexp[I-]s2
dt
(5)
(6)
k[HIO3]0[HClO4]0
(1 + k*[HIO3]0[HClO4]0)2
kexp
)
1/2
[HIO3]0[HClO4]0
1
k1/2
k*
k1/2
)
+
[HIO3]0[HClO4]0
(7)
(
)
(
)
kexp
Figure 4 shows the data in Figures 2 and 3 plotted according to

Accumulation of H2I2O3 in Reaction of I- with IO3
J. Phys. Chem. A, Vol. 104, No. 22, 2000 5271
-
1/2
Figure 4. Plot of ([HIO3]0[HClO4]0/kexp
)
× 104 (M3 s)1/2 vs ([HIO3]0-
Figure 2. Saturation of kexp/104 (M-1 s-1) with increasing ([HIO3]0-
[HClO4]0) × 103 (M2) according to eq 7. Slope ) k*/k1/2 ) (7.8 ( 0.4)
[HClO4]0) × 103 (M2) over the range 0.000 448 M e [KIO3]0
e
× 10-2 (s/M)1/2 ) (K9K10/k11)1/2 and intercept ) 1/k1/2 ) (8.0 ( 0.2) ×
0.004 48 M) in five series of experiments at constant [HClO4]0 of (b)
0.0968 M, (2) 0.199 M, (9) 0.289 M, (() 0.393 M, and (1) 0.489 M.
10-5 (M3 s)1/2 ) (1/K9K10k11)1/2
.
at constant [IO3-]0 (0.000 448-0.004 48 M) with [HClO4]0
varying and about 20 times higher (0.0968-0.489 M) than
[IO3-]0. Thus, the ionic strength varies considerably within each
of these series of experiments. Even though the data from
Figures 2 and 3 fit together well in Figure 4, this varying ionic
strength undoubtedly introduces some error into the inferred
values of k and k*.
Mechanistic Interpretation. The mechanism below is in
keeping with previous work, its recent analysis by Schmitz,14c
and with eq 5. The species H2I2O3 is suggested to be similar in
Figure 3. Saturation of kexp/104 (M-1 s-1) with increasing ([HIO3]0-
[HClO4]0) × 103 (M2) over the range 0.0968 M e [HClO4]0 e 0.489
M in five series of experiments at constant [KIO3]0 of (b) 0.000 448
M, (2) 0.000 896 M, (9) 0.001 79 M, (() 0.002 24 M, and (1) 0.004 48
M.
eq 7. The fit to eq 7 is excellent over the substantial concentra-
tion ranges cited above, and we assume eq 5 describes the
kinetics of reaction 1 under the condition [KIO3]0 . [KI]0/5.
The slope and intercept of the plot in Figure 4 yield k ) (1.6 (
0.1) × 108 M-3 s-1 and k* ) (9.9 ( 0.8) × 102 M-2. The
quoted precision of these quantities may be misleadingly high
because of medium effects. The data in Figure 2 are for five
series of experiments with varying [KIO3]0 (0.000 448-0.004 48
M) at constant and at least 20 times higher [HClO4]0 (0.0968-
0.489 M). Thus, the ionic strength is nearly constant within each
of these series of experiments, demonstrating that the basic
saturation phenomenon observed is not a medium effect.
However, the data in Figure 3 are for five series of experiments
2-
structure but not isoelectronic with the thiosulfate ion, S2O3
;
it is not possible for us to distinguish kinetically between the
intermediacy of H2I2O3 or its anhydride I2O2. Reaction 10 is to
be contrasted with the simple oxygen-transfer process presumed

5272 J. Phys. Chem. A, Vol. 104, No. 22, 2000
Agreda B. et al.
to be rate determining in reaction 2, reaction 14. It is suggested
Assuming reaction 11 to be rate-determining for the overall
process leads to eq 25.
d[I-]
d[I-]t d[H2I2O3]t
1
5
1
5
-
s ) -
+
) k [H I O ] [I-]
11 2 2 3 t
( )(
)
t
dt
dt
dt
(25)
Substitution of eqs 23 and 24 into eq 25 leads to eq 26 for the
kinetics of reaction 1.
that the difference between reactions 10 and 14 results from
the greater size and tendency of iodine to expand its valence
shell as compared to bromine, and the lower electronegativity
of iodine as compared to bromine. We see no kinetic evidence
that H2I2O3 rearranges and decomposes directly to HOI and
HIO2, reaction 15, under the conditions employed here, and we
d[I-]s
k11K9K10[HIO3]0[HClO4]0
(1 + K9K10[HIO3]0[HClO4]0)2
1
5
-
)
[I-]s2 (26)
dt
Equation 26 is identical in form to the experimental eq 5 and
suggests the following identities: k ) k11K9K10 ) (1.6 ( 0.1)
× 108 M-3s-1, k* ) K9K10 ) (9.9 ( 0.8) × 102 M-2, and k11
) (1.6 ( 0.2) × 105 M-1s-1. Substitution of these values of
k11 and K9K10 into eq 26 together with experimental values of
[HIO3]0 and [HClO4]0 yields the observed values of kexp to
within experimental error.
H2I2O3 f HOI + HIO2
(15)
thus ignore this process, even though it is likely important at
very low [I-].14c
Reaction 13 is suggested to be followed rapidly by the
reactions below, which are analogous to reactions 9-12.
A similar expression, eq 27, is recovered using the same
methods for the circumstance [I-]0 . [IO3
] where the
-
0
appropriate mass balance is [IO3-]s ) [IO3-]t + [H2I2O3]t +
[HIO3]t.
-
d[IO3
]
s
-
)
dt
2
k11K9K10[HClO4]0[I-]0
×
1 + Ka/[HClO4]0 + K9K10[HClO4]0 + K9K10[HClO4]0[I-]0
[IO3-]s (27)
Equation 27 reduces to eq 28 at sufficiently high [I-]. Equation
28 is in keeping with the observation of a term first order in
[I-] at high [I-]0.8,14
Followed by
-
d[IO3 ]s
-
-
) k11[I-]0[IO3 ]s
(28)
dt
A transition between first- and second-order kinetics in [I-]
as [I-] increases also can be rationalized via a rate expression
derived from the same mechanism but without assuming
accumulation of stoichiometrically significant amounts of
H2I2O3.14c One instead assumes that a very small, stoichiomet-
rically insignificant (compared to [I-]) [H2I2O3] is controlled
not solely by equilibrium 10, but instead by a steady-state
involving reactions 10, -10, and 11.14c Perturbation of the
iodine-atom mass balance may be neglected if only a very small
concentration of H2I2O3 is present. This approach leads to eq
29, which is first order in [I-] when k11 [I-] . k-10 and second
We are not able here to distinguish kinetically any details of
reactions 16-22.
Derivation of the rate expression predicted by this mechanism
proceeds as follows under the assumption that concentrations
of H2I2O3 are significant compared to the concentration of I-
accumulated in the course of the reaction. The experiments
reported here are based upon spectrophotometric measurement
of the concentration of I2 present at time t, i.e., [I2]t. The
stoichiometric concentration of iodide ion, [I-]s, remaining at
time t then is given by [I-]s ) 5/3([I2]∞ - [I2]t). This iodide ion
is distributed between free I- and H2I2O3 according to the mass
balance, [I-]s ) [I-]t + [H2I2O3]t. The assumption that reactions
9 and 10 are at equilibrium yields [H2I2O3]t ) K9K10[HIO3]0-
[HClO4]0[I-]t, and substitution of this expression into the mass
balance leads to eqs 23 and 24.
k10k11K9[HIO3][HClO4][I-]2
k-10 + k11[I-]
d[I-]
5 dt
1
-
)
(29)
order in [I-] when k11 [I-] , k-10. While it is very likely that
complexity resulting from the competition of reactions -10 and
11 with each other, as well as with reaction 15, is sometimes
important over the wide range of [I-] possible in this system,14c
eq 29 is not able to rationalize the saturation of kexp with ([H+]0-
[HIO3]0) that is a dominant feature of the data presented here.
We thus emphasize the formation of binuclear iodine-containing
species in this analysis.
[I-]s
[I-]t )
(23)
(24)
1 + K9K10[HIO3]0[HClO4]0
K9K10[I-]s[HIO3]0[HClO4]0
A two-term rate expression of the form of eq 30, such as
[H2I2O3]t )
considered by Liebhafsky and Roe,11 exhibits second-order
1 + K9K10[HIO3]0[HClO4]0

Accumulation of H2I2O3 in Reaction of I- with IO3
J. Phys. Chem. A, Vol. 104, No. 22, 2000 5273
-
Figure 6. Stopped-flow kinetic data, absorbance vs time, for an
experiment with [HClO4]0 ) 0.489 M; [KIO3]0 ) 0.004 48 M; and
[KI]0 ) 0.000 576 M.
from reactions 8-12, 15-17, and 21 using the rate constants
given in the legend of Figure 5. The values of k11 and (k9/k-9)-
(k10/k-10) ) K9K10 are set to the experimental values inferred
from Figure 4 and eq 26, with all other rate parameters set to
sufficiently large values that reactions 9-11 are rate-determin-
ing. Agreement between simulation and experiment is good,
within the experimental error of the rate parameters, in both
cases. The simulation in Figure 5 at relatively low values of
[HIO3]0 and [HClO4]0 (lower) shows only a small fraction of
I- tied up as H2I2O3, while the simulation at the highest values
of [HIO3]0 and [HClO4]0 used experimentally (upper) shows a
significant fraction of I- tied up as [H2I2O3]. Of note is that the
small induction period appearing experimentally in Figure 6 is
reproduced in the simulation (Figure 5 upper). This induction
period appears in the simulation because of the time required
for the formation of the intermediate H2I2O3. Its appearance in
both experiment and simulation is further evidence in support
of the assumption of accumulation of stoichiometrically sig-
nificant quantities of it in the Dushman reaction. Assuming that
the proton transfer in reaction 10 is very rapid suggests that the
value of (k10 + k-10) ≈ 2.5 × 105 M-1 s-1 assumed in the
simulation accurately reflects the time scale of formation of
H2I2O3 in reactions 9 and 10.
Figure 5. Simulation of the experiments in Figures 1 and 6 by
numerical integration of the set of differential equations resulting from
application of the law of mass action to reactions 8-12, 15-17, and
21. The assumed rate parameters are k8 ) 1.0 × 1010 M-1 s-1, k-8
)
3.2 × 109 s-1, k9 ) 6.0 × 107 M-1 s-1, k-9 ) 3.6 × 108 s-1, k10 ) 2.5
× 105 M-1 s-1, k-10 ) 42.2 s-1, k11 ) 1.6 × 105 M-1 s-1, k12 ) 1.0 ×
108 s-1, k16 ) 1.0 × 108 M-1 s-1, k-16 ) 1.0 × 103 s-1, k17 ) 1.0 ×
109 M-1 s-1, k-17 ) 1.0 × 105 s-1, k18 ) 1.0 × 106 M-1 s-1, k20 ) 1.0
× 109 M-1 s-1, k-20 ) 1.0 × 109 s-1, k21 ) 1.0 × 1010 M-2 s-1, k-21
) 100 s-1. (Lower) [I2] × 105 (M) (s) and [H2I2O3] × 105 (M) (---)
vs time (s). [HClO4]0 ) 0.289 M, [KIO3]0 ) 0.000 448 M, [HIO3]0 )
0.000 212 M, and [KI]0 ) 0.000 058 M. (Upper) [I2] × 104 (M) (s)
and [H2I2O3] × 104 (M) (---) vs time (s). [HClO4]0 ) 0.489 M, [KIO3]0
) 0.004 48 M, [HIO3]0 ) 0.002 70 M, and [KI]0 ) 0.000 057 6 M.
Discussion
Assuming the intermediacy of H2I2O3 or its anhydride I2O2
in the reaction of I- with IO3-, and that this species may
accumulate to stoichiometrically significant levels in which the
reactant present in excess complexes a substantial fraction of
the other reactant, rationalizes both our data at [IO3-]0 . [I-]0/
5 and much of the complexity of previous experimental results
under other conditions. The major experimental support for this
assumption is the appearance of saturation kinetics in [HClO4]0
and particularly [HIO3]0. The importance of H2I2O3 in the
I-/IO3- reaction, while the equivalent species H2Br2O3 does not
appear to be important in the Br-/BrO3- reaction, likely results
from the greater size and ability of iodine to expand its valence
shell as compared to bromine, and the greater electronegativity
of bromine as compared to iodine. Analogous reasoning suggests
kinetics at high [I-], and first-order kinetics at low [I-], the
d[I-]
1
5 dt
-
) ka[HIO3][H+][I-] + kb[HIO3][H+][I-]2 (30)
opposite of eqs 27 and 29 and to what is observed both here
and in previous work. Thus, we do not consider here mechanistic
paths involving two independent paths from reactants to
products, such as is implied by eq 30.
Figure 5 shows simulated values of [I2] and [H2I2O3] for the
experiments displayed in Figures 1 and 6 calculated by
numerical integration of the mass-action rate equations resulting

5274 J. Phys. Chem. A, Vol. 104, No. 22, 2000
Agreda B. et al.
that H2I2O2 or its anhydride I2O likely is important in the
reaction of HIO2 with I-.
from the stability of the species H2ClBrO3. Some work has been
done on catalysis of the I-/IO3- reaction by Br- and Cl-,14c,30-32
but it is difficult to interpret this work within the context of the
ideas presented here.
The [HClO4]0 and [HIO3]0 dependence of the I-/IO3- reaction
may be more complex than is indicated here and deserves further
investigation. The complexity of iodic acid solutions has been
considered by Furuichi and Liebhafsky,27 who also found a form
of saturation kinetics in the I-/IO3- reaction and discussed the
existence of H2I2O3. Indeed, while the data presented here are
best fitted by eq 5, the severe saturation of kexp seen at the
highest values of ([HIO3]0[H+]0) in Figure 2 may contain a
contribution from
Acknowledgment. This work is supported by Universidad
Nacional de Colombia, The University of Montana, and the
National Science Foundation under grant OSR-9350546. R.J.F.
is pleased to acknowledge very useful discussions with Professor
Guy Schmitz.
References and Notes
HIO3 + IO3- + H+ a I2O5 + H2O
(31)
(1) Dushman, S. J. Phys. Chem. 1904, 8, 453.
(2) Abel, E.; Stadler, F. Z. Phys. Chem. 1926, 122, 49.
(3) Abel, E.; Hilferding, K. Z. Phys. Chem. 1928, 136, 186.
(4) Bray, W. C. J. Am. Chem. Soc. 1921, 43, 1262.
(5) Noyes, R. M.; Sharma, K. J. Am. Chem. Soc. 1975, 97, 202.
(6) Liebhafsky, H. A.; Wu, L. S. J. Am. Chem. Soc. 1974, 96, 7180.
(7) Treindl, L.; Noyes, R. M. J. Phys. Chem. 1993, 97, 11354.
(8) Barton, A. F. M.; Wright, G. A. J. Chem. Soc. A 1968, 2096.
(9) Schildcrout, S.; Fortunato, F. A. J. Phys. Chem. 1975, 79, 31.
(10) Papsin, G.; Hanna, A.; Showalter, K. J. Phys. Chem. 1981, 85,
2575.
(11) Liebhafsky, H. A.; Roe, G. M. Int. J. Chem. Kinet. 1979, 11, 693.
(12) Rabai, G.; Kaminaga, A.; Hanzaki, I. J. Phys. Chem. 1995, 99,
9795.
(13) Mambo, E.; Simoyi, R. J. Phys. Chem. 1993, 97, 13662.
(14) (a) Schmitz, G. J. Chim. Phys. 1987, 84, 957. (b) Schmitz, G. J.
Chim. Phys. 1991, 88, 15. (c) Schmitz, G. Phys. Chem. Chem. Phys. 1999,
1, 1909. (d) Schmitz, G. Phys. Chem. Chem. Phys. 1999, 1, 4605.
(15) Urbansky, E. T.; Cooper, B. T.; Margerum, D. Inorg. Chem. 1997,
36, 1338.
We assume that I2O5 is less reactive with I- than is HIO3. This
suggestion is in keeping with the major implication of the data
reported here that stoichiometrically significant quantities of
binuclear, iodine-containing species accumulate in this system.
An induction period is observed in some of our experiments,
apparently related to the time required for formation of H2I2O3.
Furthermore, at some wavelengths a transient absorbance is
observed, perhaps related to H2I2O3. Considerably more work
is required to clarify these observations. The I-/IO3 reaction
-
also should be investigated by the stopped-flow kinetic methods
for [IO3-]0 , [I-]0/5.
The mixed halogen reactions, I-/BrO3- and Br-/IO3-, are a
significant test of the suggestion that H2I2O3 forms readily and
is quite stable, while H2Br2O3 does not. It is expected that
(16) Field, R. J.; Fo¨rsterling, H.-D. J. Phys. Chem. 1986, 90, 5400.
(17) Myers, O. E. J. Chem. Phys. 1958, 28, 1027.
(18) Matsuzaki, I.; Nakajima, T.; Liebhafsky, H. A. Chemistry Lett.
Chem. Soc. Jpn. 1974, 1463.
(19) Matsuzaki, I.; Nakajima, T.; Liebhafsky, H. A. Faraday Symp.
Chem. Soc. 1975, 9, 55.
(20) Liebhafsky, H. A.; McGavock, W. C.; Reyes, R. J.; Wu, L. S. J.
Am. Chem. Soc. 1978, 100, 87.
(21) Barton, A. F. M.; Cheong, H. N.; Smidt, R. E. J. Chem. Soc.,
Faraday Trans. 1976, 72, 568.
(22) Kolar-Anic´, L. Z.; Misˇljenovic´, D. M.; Stanisavljev, D. R.; Anic´,
S. R. J. Phys. Chem. 1990, 94, 8144.
-
will occur in the Br-/IO3 reaction, leading under some
conditions to kinetics second order in [Br-], while
(23) Kolar-Anic´, L.; Schmitz, G. J. Chem. Soc., Faraday Trans. 1992,
88, 2343.
(24) (a) Kolar-Anic´, L.; CÄ upic, Zˇ.; Anic´, S.; Schmitz, G. J. Chem. Soc.,
Faraday Trans. 1997, 93, 2147 (b) Cupic, Zˇ; Kolar-Anic´, L. J. Chem. Phys.
1999, 110, 3951.
(25) Strong, L. E.; Pethybridge, A. D. 1987, 16, 891.
(26) Woodbury, G. Physical Chemistry, Brooks/Cole: Pacific Grove,
CA, 1997; p 258.
(27) Furuichi, R.; Liebhafsky, H. A. Bull. Chem. Soc. Jpn. 1975, 48,
745.
(28) (a) Barton, A. F. M.; Wright, G. A. J. Chem. Soc. A 1968, 1747.
(b) Barton, A. F. M.; Loo, B. H. J. Chem. Soc. A 1971, 3032.
(29) Sigalla, J. J. Chim. Phys. 1958, 55, 758.
(30) Beran, P.; Bruckenstein, S. J. Phys. Chem. 1968, 72, 3630.
(31) Nikolelis, D. P.; Karayannis, M. I.; Hadjiioannou, T. P. Anal. Chem.
Acta 1977, 94, 415.
will not appear in the I-/BrO3 reaction, which will occur
-
instead via an oxygen-transfer mechanism as in the Br-/BrO3
-
reaction, leading to kinetics first order in [I-]. The kinetics of
-
the I-/BrO3 reaction has been investigated21,28 and is indeed
first order in [I-], but we know of no definitive kinetic
-
investigation of the Br-/IO3 reaction. This suggestion is
supported by the observation29 of a second-order term in the
kinetics of the oxidation of Cl- by BrO3-, presumably resulting
(32) Gaspar, V.; Beck, M. T. J. Phys. Chem. 1986, 90, 6303.