Oxyhalogen-sulfur chemistry - Kinetics and mechanism of oxidation of methionine by aqueous iodine and acidified iodate

Article abstract of DOI:10.1139/V09-038

The oxidation of methionine (Met) by acidic iodate and aqueous iodine was studied. Though the reaction is a simple two-electron oxidation to give methionine sulfoxide (Met-S=O), the dynamics of the reaction are, however, very complex, characterized by clock reaction characteristics and transient formation of iodine. In excess methionine conditions, the stoichiometry of the reaction was deduced to be IO₃^- + 3Met → I^- + 3Met-S=O. In excess iodate, the iodide product reacts with iodate to give a final product of molecular iodine and a 2:5 stoichiometry: 2IO₃ ^- + 5Met + 2H⁺ → I₂ + 5Met-S=O + H ₂O. The direct reaction of iodine and methionine is slow and mildly autoinhibitory, which explains the transient formation of iodine, even in conditions of excess methionine in which iodine is not a final product. The whole reaction scheme could be simulated by a simple network of 11 reactions.

Full text of DOI:10.1139/V09-038

689
Oxyhalogen–sulfur chemistry — Kinetics and
mechanism of oxidation of methionine by aqueous
iodine and acidified iodate
Edward Chikwana, Bradley Davis, Moshood K. Morakinyo, and Reuben H. Simoyi
Abstract: The oxidation of methionine (Met) by acidic iodate and aqueous iodine was studied. Though the reaction is a sim-
ple two-electron oxidation to give methionine sulfoxide (Met–S=O), the dynamics of the reaction are, however, very com-
plex, characterized by clock reaction characteristics and transient formation of iodine. In excess methionine conditions, the
stoichiometry of the reaction was deduced to be IO3– + 3Met ? I– + 3Met–S=O. In excess iodate, the iodide product reacts
with iodate to give a final product of molecular iodine and a 2:5 stoichiometry: 2IO3– + 5Met + 2H+ ? I2 + 5Met–S=O +
H2O. The direct reaction of iodine and methionine is slow and mildly autoinhibitory, which explains the transient formation
of iodine, even in conditions of excess methionine in which iodine is not a final product. The whole reaction scheme could
be simulated by a simple network of 11 reactions.
Key words: methionine, organosulfur, iodine, antioxidant, nonlinear dynamics, methionine sulfoxide.
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Resume : On a etudie la reaction d’oxydation de la methionine (Met) par l’ion iodate, en milieu acide et en presence
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d’iode aqueux. Meme si la reaction d’une simple oxydation a deux electrons pour conduire au sulfoxyde de methionine
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(Met–S=O), la dynamique de la reaction est toutefois fort complexe ayant les caracteristiques d’une reaction horloge et
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comportant la formation transitoire d’iode. Dans des conditions ou la methionine est en exces, on a deduit que la stoechio-
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metrie de la reaction est IO3 + 3Met ? I + 3Met–S=O. Dans des conditions ou l’iodate est en exces, l’iodure produit
reagit avec l’iodate pour donner de l’iode moleculaire comme produit final avec une stoechiometrie 2:5: 2IO3 + 5Met +
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+
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2H ? I2 + 5Met–S=O + H2O. La reaction directe de l’iode avec la methionine est tres lente et donne lieu a une legere
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autoinhibition qui explique la formation transitoire d’iode, meme dans des conditions d’exces de methionine dans lesquel-
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les l’iode n’est pas un produit final. L’ensemble des reactions peut etre simule par un reseau simple de onze reactions.
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Mots-cles : methionine, organosulfure, iode, antioxydant, dynamique non lineaire, sulfoxyde de methionine.
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[Traduit par la Redaction]
pecially in combination with chemotherapy.4 On the other
hand, high levels of methionine in the physiological system
are associated with several metabolic disorders such as ho-
mocystinuria, which is characterized by increased levels of
homocysteine in the serum.5 Homocystine is one of the
products of methionine metabolism and its accumulation
has been linked to atherosclerosis and arteriosclerosis.1,5
Atherosclerosis is characterized by the development of le-
sions within and around blood vessels.5,6 While atheroscle-
rosis provides the plaque that narrows the artery,
arteriosclerosis stiffens the arteries so that they cannot ex-
pand with each heart beat to compensate for the blockage
caused by plaque formation.6,7 In most cases of myocardial
infarction, both atherosclerosis and arteriosclerosis have
been shown to be present.
Introduction
Methionine is an essential amino acid, obtained only from
dietary sources and found mainly in animal protein. In addi-
tion to its role as a precursor in protein synthesis, L-methio-
nine participates in a wide range of biochemical reactions,
including the production of S-adenosyl methionine, homo-
cysteine, L-cysteine, taurine, and sulfate. It is believed to be
the major methyl donor for methylation of DNA, RNA, pro-
teins, and other molecules.1,2 Methionine is also thought to
keep fat from building up in the liver, and it is often in-
cluded in liver-detoxifying products called lipotropic combi-
nations.3 These formulations are believed to accelerate the
flow of bile and cell-damaging toxins away from the liver.
Depletion of methionine is strongly linked to inhibited
growth, apoptotic death, and necrosis of cancerous cells, es-
It has been proposed that methionine residues in proteins
serve as an antioxidant defense system for the protection of
proteins from oxidation under conditions of oxidative
stress.8,9 Various biological oxidants, such as hydrogen per-
oxide, hydroxyl radicals, ozone, peroxynitrite, and hypo-
chlorite, have been shown to mediate the oxidation of
methionine.9,10 Methionine has been shown to be effective
in radical scavenging processes, such as those involving re-
active nitrogen species, to form methionine sulfoxide and
ethylene.9 In the physiological environment, oxidation of
methionine residues to methionine sulfoxide is easily re-
Received 24 October 2008. Accepted 24 December 2008.
Published on the NRC Research Press Web site at
canjchem.nrc.ca on 7 May 2009.
E. Chikwana,1 B. Davis, M.K. Morakinyo, and R.H. Simoyi.2
Department of Chemistry, Portland State University, Portland,
OR 97207-0751, USA.
1Present address: Department of Chemistry, Franklin College,
Franklin, IN 46131-2623, USA.
2Corresponding author (e-mail: rsimoyi@pdx.edu).
Can. J. Chem. 87: 689–697 (2009)
doi:10.1139/V09-038
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Can. J. Chem. Vol. 87, 2009
Fig. 1. Titration results for the oxidation of methionine in ex-
cess iodate conditions. The x-intercept strongly suggest a 1:3
(iodate:methionine) stoichiometric ratio. [Met]0 = 3.0 ꢀ 10–3 mol/L,
[H+]0 = 6 ꢀ 10–3 mol/L.
Fig. 3. Iodate variation shows transient iodine formation below the
stoichiometry (traces a–c). In excess iodate conditions, the final io-
dine amount reaches a maximum determined by the methionine
concentration as shown by traces e–h. The relation between iodate
concentration and induction time is also apparent. [Met]0 =
0.005 mol/L, [H+]0 = 0.02 mol/L, [IO3 ]0 = (a) 0.001, (b) 0.00125,
–
(c) 0.0015, (d) 0.00175, (e) 0.002, (f) 0.0025, (g) 0.003, (h)
0.004 mol/L.
Fig. 2. (A) NMR spectrum for the product of the DL-methionine
oxidation by acidic iodate/iodine solution showing the S-methyl
peak at 2.79 ppm. (B) Spectrum of DL-methionine shows the peak
initially at 2.12 ppm. Spectra for reagent grade methionine sulfox-
ide (C) and methionine sulfone (D) shows the shifting in the S-
methyl group as expected from the oxidation of the sulfur center.
The product solution in this case shows strong similarity to the
spectrum of the sulfoxide as predicted from the stoichiometric de-
termination.
system, such as the oxygen radical species, hydrogen perox-
ide, hypobromous acid, hypochlorous acid, and iodine, can
go a long way in providing data on the depletion of exces-
sive dietary methionine in tissues that, unlike the liver, do
not contain the complete methionine cycle.
Methionine has also been shown to inhibit in vivo thyroid
peroxidase (TPO) iodide oxidation and iodination activ-
ities.11,14 Its reaction with iodine is expected to play a major
role in its interaction with the thyroid enzymes. Several or-
ganosulfur compounds are well-known goitrogenics whose
actions are based on their ability to abstract the active iodine
cation from the thyroid, thereby reducing its hyperactiv-
ity.15–19 We report, in this manuscript, on a detailed kinetics
and mechanistic study of the oxidation by iodate/iodine of
DL-methionine. This simple oxidation reaction was surpris-
ingly complex, characterized by nonlinear exotic kinetics
behavior.
Experimental
Materials
Iodine, potassium iodide (Sigma-Aldrich), sodium
perchlorate (98%) (Acros Organics), potassium iodate, DL-
methionine (99%), DL-methionine sulfone (99%), DL-
methionine sulfoxide, perchloric acid (72%), soluble starch,
sodium thiosulfate, and hydrochloric acid (Fisher Scientific)
were used without further purification. The concentration of
iodine was determined by standardization against thiosulfate
with starch as the indicator. Spectrophotometry was also uti-
lized by measuring iodine absorbance at its isosbestic point
with triiodide at 460 nm where the extinction coefficient had
been deduced to be 770 (mol/L)–1 cm–1. This standardization
was carried out before each series of kinetics experiments
due to the volatile nature of iodine. Methionine solutions
were prepared just before use and not kept for more than
versed by the action of the enzyme methionine sulfoxide re-
ductase.9 Even though methionine has three coordinating
centers (O, N, and S), studies have shown that the sulfur
center is the most susceptible to oxidative attack.9,11–13
The key to understanding the physiological role of me-
thionine requires an understanding of its specialized func-
tion, its reactive intermediates, and oxidation products.
Even though most of the methionine is believed to be me-
tabolized in the liver via the transmethylation and transsulfu-
ration pathways,1 alternative pathways cannot be excluded.
Its interaction with different oxidants in the physiological
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Chikwana et al.
691
Fig. 5. Absorbance traces for methionine variation in its oxidation
Fig. 4. (a) The effect of varying acid concentrations for reactions run
in excess iodate. Acid is not a reactant in the reaction but strongly
catalyzes the reaction by reducing the induction period. [Met]0 =
0.003 mol/L, [IO3 ]0 = 0.03 mol/L, [H+]0 = (a) 0.003, (b) 0.004,
by iodate. Initially, the final iodine absorbances increase (a–d) but
as the stoichiometric point is approached they start to decrease and
show transient iodine formation (e). [IO3 ]0 = 0.003 mol/L, [H+]0 =
–
–
(c) 0.0045, (d) 0.005, (e) 0.006, (f) 0.007, (g) 0.008 mol/L. (b) The
effect of acid variation in excess methionine. The effect of increas-
ing acid can be seen in the form of a shorter induction period and a
higher maximum transient iodine concentration. [Met]0 =
0.0045 mol/L, [Met]0 = (a) 0.003, (b) 0.004, (c) 0.005, (d) 0.006,
(e) 0.007 mol/L.
0.005 mol/L, [IO3 ]0 = 0.001 mol/L, [H+]0 = (a) 0.0125, (b) 0.015,
–
(c) 0.0175, (d) 0.02, (e) 0.025, (f) 0.03 mol/L.
Fig. 6. Iodide effect on the iodate oxidation of methionine. The ef-
fect of iodide can be seen on the reduced induction period and in-
creased final absorbance of iodine as iodide concentration is
increased. [Met]0 = 0.003 mol/L, [H+] = 0.0045 mol/L, [IO3 ] =
–
0.0275 mol/L, [I–] = (a) no iodide, (b) 6 ꢀ 10–6, (c) 9 ꢀ 10–6,
(d) 5 ꢀ 10–5, (e) 1 ꢀ 10–4 mol/L.
24 h. All solutions were prepared using distilled water from
a Barnstead Sybron Corporation water purification unit. In-
ductively coupled plasma mass spectrometry (ICPMS) was
used to show that our reaction medium did not contain
enough copper ions to affect the overall reaction kinetics
and mechanism.20
trophotometer as well as on a Hi-Tech Scientific SF-DX2
double-mixing stopped-flow spectrophotometer.
Stoichiometric determinations
The stoichiometry for the methionine–IO3 reaction was
–
Methods
All experiments were carried out at 25.0
0.1 8C and
determined both in excess iodate and excess methionine
conditions. In excess iodate the total excess oxidizing power
was determined by titration. Excess acidified iodide was
added to the reaction solution and the released iodine was
titrated against standard thiosulfate. Spectrophotometric
measurements were also used to determine the amount of io-
dine formed in excess iodate by its absorbance at 460 nm. In
the I2–methionine reaction, the stoichiometry was deter-
mined by titrating standardized iodine solution from a bu-
with a constant ionic strength of 1.0 mol/L (NaClO4). Me-
thionine, sodium perchlorate, and perchloric acid solutions
were mixed in one vessel and iodate (or iodine) solutions in
another. Kinetics measurements for the slower reactions and
spectrophotometric determinations were performed on a Per-
kinElmer Lambda 25 UV–vis Spectrophotometer. The faster
reactions, especially those involving iodine oxidations were
monitored on a Hi-Tech Scientific SF-61 stopped-flow spec-
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Can. J. Chem. Vol. 87, 2009
Fig. 7. (a) Traces showing variation of Methionine concentration
with constant iodine concentration. As the methionine is increased,
the reaction takes a much shorter time to reach completion. [I2]0 =
0.0003 mol/L; [Met]0 = (a) 0.01, (b) 0.015, (c) 0.02, (d) 0.025,
(e) 0.03, (f) 0.04, (g) 0.045, (h) 0.05 mol/L. (b) The data shown in
Fig. 7a shows a linear dependence between initial rate of reaction
and methionine concentration.
Fig. 8. (a) Effect of deliberately added iodide (product of the reac-
tion) on the rate of reaction. [Met] = 5.0 ꢀ 10–3 mol/L; [I2] = 2.2 ꢀ
10–4 mol/L ; [I–] = (a) no added iodide; (b) 5.0 ꢀ 10–5; (c) 1.0 ꢀ 10–4;
(d) 2.0 ꢀ 10–4; (e) 3.0 ꢀ 10–4; (f) 4.0 ꢀ 10–4; (g) 5.0 ꢀ 10–4 mol/L.
(b) The inverse plot of initial rate and iodide concentrations for the
data in Fig. 8a. Plot assumes I3– is inert. The inability of these data to
reproduce the rate constant and Keq indicates that the I3– is not inert.
rette into a solution of methionine of known strength. The
end point, which was enhanced by the starch indicator, was
detected as the point where the blue-black color lingers. In
excess iodine conditions, spectrophotometric determinations
were used to complement the titrimetric techniques.
axis.21 The data in Fig. 1, which were derived using a constant
concentration of 0.003 mol/L methionine, show that there is
no excess oxidizing power when [IO3 ]0 = 0.001 mol/L, which
–
simplifies to a ratio of 1:3. In excess methionine, the stoichi-
ometry for the reaction involved 1 mol of iodate reacting with
3 moles of methionine. This suggests a two-electron oxidation
of the sulfur center, with methionine sulfoxide as the most
likely product. Titrimetric and spectrometric techniques were
both used to verify the following stoichiometry
Results
Stoichiometry
Solutions with varying ratios of iodate to methionine were
prepared and their excess oxidizing power determined by io-
dometric titrations. Excess oxidant was acidified and mixed
with excess iodide to generate iodine according to the Dush-
man reaction, eq. [R1].
½R2ꢁ
½R1ꢁ
IO3ꢂ þ 5Iꢂ þ 6Hþ ! 3I2 þ 3H2O
The iodine was titrated against standard sodium thiosulfate
and a plot of thiosulfate titer required versus initial iodate con-
centration was made and is shown in Fig. 1. The number of
moles of iodate that were required to react with methionine
without formation of iodine (from excess iodate) can be deter-
mined from extrapolation to the methionine concentration
In stoichiometric excess of iodate, in addition to MetS=O,
iodine is also produced with an overall reaction stoichiome-
try of
Published by NRC Research Press

Chikwana et al.
693
Scheme 1. Mechanism of reaction of methionine with iodic acid.
oxidation resulting in cleavage to the C–S bond. Proton NMR
spectroscopy was used to confirm the products as predicted
from the stoichiometry. In Fig. 2, spectrum A is of the prod-
uct of reaction between acidic iodate and methionine. Spec-
trum B is pure DL-methionine, which shows the expected
strong singlet for the S-methyl group at 2.12 ppm. These pro-
tons appear as a singlet due to the presence of the sulfur,
which acts as a shield from the splitting effects of the other
protons on adjacent carbon atoms. Spectrum C shown is of
reagent grade DL-methionine sulfoxide, showing a downfield
shift of the methyl protons (2.79 ppm) due to the presence of
the oxygen on the sulfur center. Addition of another oxygen
to form the sulfone results in a further downfield shift of
these protons to 3.20 ppm. Spectrum D is reagent grade
methionine sulfone. Spectrum A, the reaction product, is sim-
ilar to the spectrum for the reagent grade sulfoxide (C).
Solutions to obtain spectra B–D were prepared in neutral con-
ditions while DCl was used to acidify the solution used to ob-
tain spectrum A. The difference in pH between these
solutions accounts for the small difference between spectra A
and C. The spectrum of oxidation of methionine by acidified
iodate shows that there is a small amount of sulfone formed
as a minor product, but this could not be detected iodometri-
cally due to the oxidation of iodide by the sulfone to iodine
and the sulfoxide. To further confirm the sulfoxide as the ma-
jor product, reagent grade sulfoxide was mixed with acidified
iodate solution and no further reaction was observed.
½R3ꢁ
2IO3ꢂ þ 5Met þ 2Hþ ! I2 þ 5MetꢂS¼O þ H2O
Reaction dynamics
The reaction exhibited clock behavior22 with transient io-
dine formation. Figure 3 shows both aspects of this dynami-
cal behavior. When the oxidant-to-reductant ratio,
R ¼ ½IO3ꢂꢁ0=½Metꢁ0, is less than 0.333, no iodine is expected
to be formed at the end of the reaction as predicted by stoi-
chiometry eq. [R2]. Traces a–c in Fig. 3 are derived from
reaction solutions in which R < 0.333, but, however, show
transient iodine formation, which is an indication of the
dominance of the reactions that form iodine over those that
consume it during the initial stages of the reaction. As the
reaction proceeds, iodine consumption reactions later be-
come dominant. Clock reaction characteristics, in which
there is an initial quiescent period followed by formation of
iodine, are also exhibited for all values of R, as shown by
traces a–h (Fig. 3). In stoichiometric excess of oxidant,
which gives stoichiometry eq. [R3], i.e., R ‡ 0.333, the final
iodine formation is determined by the concentration of me-
thionine. When R > 0.4, however, the amount of iodine
formed becomes invariant to further increases in iodate con-
centrations. The increase in final iodine concentration be-
tween traces d and e (Fig. 3) is justified based on the fact
that stoichiometry eq. [R3] is a step wise linear combination
of eqs. [R1] and [R2].
Essentially, stoichiometry eq. [R3] is a linear combination
of stoichiometry eqs. [R1] and [R2]: [R1] + 5[R2]. Iodide
produced in stoichiometry eq. [R1] is consumed by the
excess iodate to produce iodine, and the end of reaction is
determined by a complete consumption of iodide for as
long as there is still excess iodate to fuel the Dushman reac-
tion, eq. [R1]. Initially there is an increase in the amount of
iodine produced as the iodate to substrate ratio,
R ¼ ½IO3ꢂꢁ0=½Metꢁ0, is increased from 0.333 (stoichiometry
eq. [R2]), until a saturation point is reached where the me-
thionine becomes limiting at a ratio of 0.400. In this short
range of ratios, the final iodine amount was reproducibly de-
termined spectrophotometrically to be 0.2 times the amount
of methionine used as predicted by stoichiometry eq. [R3].
The direct reaction of methionine with iodine was rapid
enough for the stoichiometry to be determined by direct ti-
tration of an iodine solution with methionine and vice-versa.
The detection of the end point was aided by the starch indi-
cator and the stoichiometry was determined to be
½R4ꢁ
Met þ I2 þ H2O ! MetꢂS¼O þ 2Iꢂ þ 2Hþ
Stoichiometry eq. [R4] could also be obtained spectropho-
tometrically in excess iodine conditions by observing the ab-
sorbance of the excess iodine at 460 nm (at the end of the
reaction).
Figure 3 also shows that as the iodate concentration is in-
creased, the rate of iodine formation increases while the qui-
escent time before iodine formation commences (induction
period), decreases.23 Formation of iodine, in all traces
(Fig. 3), is not sharp, but gradual. Thus the induction period
was estimated to be the point where rapid formation of io-
dine commenced.
Confirmation of products
The two major products that could possibly be obtained
from the oxyhalogen oxidation of the sulfur center in methio-
nine are the sulfoxide and the sulfone. Due to the lack of
BaSO4 precipitation, there was no evidence to suggest further
Figure 4a shows that not only does an increase in acid de-
crease the induction period, but it also increases the rate of
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694
Can. J. Chem. Vol. 87, 2009
Table 1. Iodate–Iodine–Methionine reaction network.
Number
M1
M2
M3
M4
M5
M6
M7
M8
Reaꢂction
Kf; kr
IO3 þ Iꢂ þ 2Hþ Ð HIO2 þ HOI
HIO2 þ Iꢂ þ Hþ Ð 2HOI
HOI þ Iꢂ þ Hþ Ð I2 þ H2O
2.8; 1.44 ꢀ 103
2.1 ꢀ 108; 90
3.1 ꢀ 1012; 2.2
8.6 ꢀ 102; 2.00
ꢂ
IO3 þ HOI þ Hþ Ð 2HIO2
I2 þ Iꢂ Ð Iꢂ3
6.2 ꢀ 109; 8.5 ꢀ 106
ꢂ
IO3 þ Hþ Ð HIO3
5.0 ꢀ 109; 1.25 ꢀ 109
HIO3 þ Met þ Hþ ! HIO2 þ MetꢂSO þ Hþ
50
HIO2 þ Met ! HOI þ MetꢂSO
125
165
72
M9
M10
M11
HOI þ Met ! MetꢂSO þ Iꢂ þ Hþ
I2 þ Met þ H2O ! MetꢂSO þ 2Hþ þ 2Iꢂ
ꢂ
I3 þ Met þ H2O ! MetꢂSO þ 2Hþ þ 3Iꢂ
16
Fig. 9. Computer modeling: traces showing the simulated results
concentration. As R is reduced from a–d (Fig. 5), the final
iodine amount increases to a final concentration that is de-
termined by stoichiometry eq. [R3]. As reaction stoichiome-
try eq. [R2] is approached, the monotonic formation of
iodine gives way to transient formation of iodine as shown
in trace e (Fig. 5).
(circles) and the experimental data (solid line) for trace g shown in
Fig. 4a. [Met]0 = 0.003 mol/L, [IO3 ]0 = 0.03 mol/L, [H+] =
–
0.008 mol/L.
In excess oxidant conditions, the effect of iodide can easily
be seen by the increase in final iodine concentration. This is
expected in conditions where R > 0.4, in which the iodide
ions initially added to the reaction mixture combine with
those produced by the reaction in eq. [R2] to fuel the Dush-
man reaction,24,26 eq. [R1]. This is clearly evident in Fig. 6
where an increase in iodide concentrations increases the final
iodine concentration as well as reducing the induction time.
The direct reaction of iodine and methionine plays a cru-
cial role in this mechanism. The rate at which iodine oxi-
dizes methionine determines whether methionine can
coexist with iodine. If this reaction is fast, then final iodine
formation would mean that all the methionine has been con-
sumed. Figure 7a shows the iodine consumption of methio-
nine in high excess of methionine. Despite this high excess,
no pseudo-first-order kinetics are observed, indicating that
this is not a simple electrophilic attack of iodine on the thi-
oether group. There is a linear dependence between the ini-
tial rate of reaction with both methionine and iodine
concentrations. Figure 7b shows the methionine dependence.
Figure 8a shows that iodide is inhibitory to the oxidation of
methionine by iodine. Figure 8b shows that there is an in-
verse dependence relationship between the initial rate of re-
action and iodide concentrations. Iodide is a product of the
iodine–methionine reaction, and thus its effect is autoinhibi-
tory. This explains why no pseudo-first-order kinetics were
observed in the data shown in Fig. 7a.
formation of iodine. The catalytic effect of acid on the Dush-
man24–27 reaction (eq. [R1]) is responsible for the increase in
rate of formation of iodine in conditions where R > 0.4.
Since the final amount of iodine does not change with an in-
crease in acid, acid influences only the overall rate of forma-
tion of iodine, which in turn affects the induction period. The
effect of acid for low R values (R < 0.33), however, is more
complex as shown in Fig. 4b. Figure 4b shows traces in ex-
cess methionine conditions where no final iodine accumula-
tion is expected. As the acid is increased from a to f
(Fig. 4b), the induction period is shortened while the maxi-
mum amount of transient iodine is increased. This indicates
acid catalysis in formation of iodine, resulting in rapid accu-
mulation of iodine and shorter induction periods as acid con-
centrations are increased. Higher acid traces in Fig. 4b also
show a narrower ‘‘excursion’’ indicating that an increase in
acid also increases the rate of consumption of iodine. Given
that the general shape of the traces is not symmetrical, it
would also appear that acid catalyzes the formation of iodine
to a greater extent than it catalyzes its consumption.
Mechanism
The bulk of iodate-based oxidations are borne by the reac-
tive oxyiodine species HIO2, IO2, and HOI as well as molec-
ular iodine, I2. These species can be generated in acidic
iodate conditions, and the rate of iodate oxidations are usu-
ally controlled by the rate at which these reactive species are
generated.28,29 There is some debate as to the exact initiation
sequence of iodate oxidations, but a nucleophilic attack on a
protonated iodate species (see reaction eq. [R7]) appears to
be the most widely accepted initiation to produce iodous
and hypoiodous acids.
The oxidant-to-reductant ratio, R, plays an important role
in determining the dynamics of the reaction as can be seen
in Fig. 5, which shows the effect of progressively increasing
the methionine concentration at constant iodate and acid
Published by NRC Research Press

Chikwana et al.
695
Scheme 2. Schematic representation of all possible reactions, intermediates, and products in the oxidation of methionine by acidic iodate
and aqueous iodine.
½R5ꢁ
IO3ꢂ þ Hþ Ð IO3H
kinetics. The expected iodide concentrations, which exist
in these iodate solutions, should be approximately 5.0 ꢀ
10–7 mol/L. Assuming eqs. [R9] and [R10] are relatively
rapid, one can utilize reaction eq. [R8] to determine the
rate of formation of iodide and subsequently the rate of
oxidation of methionine. A simple integration gives a sig-
moidal curve for iodide generation. Literature values give
the composite forward rate constant of reaction eq. [R8]
as 2.2 (mol/L)–3 s–1. Using these values, one generates a
very slow rate of iodide formation and subsequent rate of
reaction, giving reaction times in the order of hours instead
of the observed 2–10 min. This suggests that iodide forma-
tion is not solely through the composite reaction eq. [R11].
The next flaw arises from the effect of iodide observed in
Fig. 6. With cubic autocatalysis, as profiled in reaction eq.
[R11], addition of minute amounts of iodide should result
in a greatly enhanced rate of reaction. Figure 6 shows
that, while iodide is catalytic, it is not as strong as would
have been expected from an autocatalyst.
Investigators of the Briggs–Rauscher30,31 and Bray–
Liebhafsky25,32 reactions have suggested that this nucleo-
philic species is the trace amount of iodide ions that are
contained in normal iodate solutions.
½R6ꢁ
IO3H þ Hþ Ð þIOðOHÞ2
½R7ꢁ
þIOðOHÞ2 þ Iꢂ Ð IꢂIOðOHÞ2
The H2I2O3 species formed in eq. [R7] can then break
down in a rate-determining step to produce the reactive spe-
cies, HIO2 and HOI.29
½R8ꢁ
IꢂIOðOHÞ2 ! HIO2 þ HOI
HOI is extremely reactive as an oxidizing species, and
will rapidly react with the substrate, methionine, to produce
iodide and the oxidation product, methionine sulfoxide.
It appears then, that initial formation of iodide should be
derived from the substrate itself acting as a nucleophile on
either the iodic acid or the protonated iodic acid. Both sce-
narios would produce kinetically indistinguishable reaction
profiles and mechanisms. In Scheme 1, methionine is repre-
sented as R1SR2.
The iodous acid produced in Scheme 1 can be further re-
duced successively by methionine to hypoiodous acid and
iodide. Reaction eq. [R9] is a rapid oxygen transfer reaction,
and thus, the overall rate of reaction should be the rate of
formation of HOI.
½R9ꢁ
HOI þ :SR1R2 ! R1R2S¼O þ Hþ þ Iꢂ
The iodide produced in eq. [R9] feeds back into reaction
eq. [R7] to facilitate further oxidation of the reactive spe-
cies. While we may accept iodous acid as a viable oxidant,
it, however, rapidly reacts with iodide to produce more hy-
poiodous acid:
½R10ꢁ
HIO2 þ Hþ þ Iꢂ Ð 2HOI
Addition of the reactions leading to the oxidation of me-
thionine in the initial stages shows that the reaction’s initial
activity is the build-up of iodide, which is autocatalytically
produced in these early stages of the reaction of eqs. [R5] +
[R6] + [R7] + [R8] + 3[R9] + [R10];
Accumulation of iodide will then transfer the rate of reac-
tion to standard Dushman reaction kinetics with the compo-
site rate-determining step eq. [R12]:
½R12ꢁ
IO3ꢂ þ Iꢂ þ 2Hþ Ð HIO2 þ HOI
½R11ꢁ
IO3ꢂ þ 2Iꢂ þ 3MetS ! 3Iꢂ þ 3MetSO
Iodine is formed from the reverse of its hydrolysis reaction:
The iodide material balance for the initial stages suggests
a mechanism that is based on cubic autocatalysis in iodide.
There are, however, two major flaws with the eqs. [R5]–
[R10] sequence. The first flaw is from pure mass action
½R13ꢁ
HOI þ Iꢂ þ Hþ Ð I2ðaqÞ þ H2O
Through relaxation spectroscopy, reaction eq. [R13] is
known to be very rapid and acid-catalyzed.33 Basic condi-
Published by NRC Research Press

696
Can. J. Chem. Vol. 87, 2009
½R17ꢁ
R1R2S: þ I3ꢂ þ H2O ! R1R2S¼O þ 3Iꢂ þ 2Hþ
tions catalyze and support the hydrolysis reaction. Our ex-
periments, however, were carried out in acidic conditions
since iodate is essentially inert in basic environments. A
summary of all possible reactions, intermediates, and prod-
ucts in the oxidation of methionine by acidic iodate and its
oxidative derivatives is shown in Scheme 2.
Combining reaction eqs. [R14] and [R17] gives a rate law
of the form
½Metꢁ0½I2ꢁ0
½2ꢁ
Rate ¼
½k14 þ k17Keq½Iꢂꢁꢁ
1 þ Keq½Iꢂꢁ
The methionine–iodine reaction
In the initial stages, before iodide accumulates, rate law
The direct reaction of methionine with aqueous iodine is
very important in determining the overall global dynamics
observed with respect to iodine formation. Transient iodine
formation, as observed in Fig. 4b, is derived from the cou-
pling of the reaction that forms iodine, eq. [R13],33 with
those that consume iodine. If the methionine-iodine reaction
is extremely rapid (e.g., diffusion controlled), no transient
iodine formation would be observed and the end of the in-
duction period would be very sharp, and would signify com-
plete consumption of methionine. The nonlinear kinetics
observed indicate that the reactions involved in the forma-
tion and consumption of iodine are comparable in magnitude
and each set asserts itself based on the availability of the
relevant reagents and mass action kinetics.
eq. [2] simplifies to a simple bimolecular rate law. In the
limit of high initial iodide concentrations, in which one can
assume constant iodide concentrations, an upper limit rate
constant for k17 was derived as 16.0 (mol/L)–1 s–1.
Computer simulation
The reaction scheme was modeled using the Kintecus
software generated by James Ianni.41 The full mechanism,
which takes into account the formation and consumption of
iodine, is shown in Table 1. It is a very simple mechanism
made up of 11 reactions with the most important reactions
being M1 and M7 (Table 1), which are responsible for ini-
tiating the reaction. Since iodate oxidations only proceed at
reasonable rates in highly acidic media, the simulations
could be simplified by assuming that acid concentrations
were buffered throughout the duration of the reaction. By
using a constant acid concentration, reactions M2–M4 and
M7 (Table 1) could be reduced to the simpler bimolecular
kinetics. Reactions M1–M4 (Table 1) are the standard oxy-
iodine reactions whose rate constants are well-known and
were derived from literature.24–27 These were not altered
during the modeling exercises. Reaction M5 (Table 1) was
studied by laser-Raman techniques and the forward and re-
verse rate constants and subsequent equilibrium constant
were also derived from literature and also not varied.37,38
Reaction M6 (Table 1) is a rapid protolytic process in which
the only relevant parameter was the assumed Ka of iodic
acid. Literature reports a wide range of values for this Ka.39
The Ka value is apparently heavily dependent on the reaction
medium, especially acid concentrations.29 Elegant experi-
mental work done by Naidich and Ricci40 in 1939 estab-
lished that concentrations of iodic acid become significant
in acidic iodate solutions at a pH of less than 2.00. All our
reactions were run at pH conditions lower than this value.
The final adopted value of 0.400 was extrapolated from the
barium iodate solubility data of Naidichi and Ricci.40 The
simulations were insensitive to the adopted values of the for-
ward and reverse rate constants for M6 (Table 1) for as long
as they were both rapid and linked by Ka. More realistic dif-
fusion-controlled rate constants for M6 (Table 1) stiffened
the simulations considerably and slowed the calculations
but without affecting the resulting simulations. Reactions
M7–M9 (Table 1) represent the initiation reactions in the
absence of initially added iodide. Reactions M8 and M9
(Table 1) are not important in determining the observed re-
action dynamics for as long as they are faster than M7
(Table 1). Rate constants for reactions M10 and M11
(Table 1) were determined from this study. Effectively,
since half of the reactions’ parameters were obtained from
literature values, the only parameters that could be altered
for best fit to the data were Ka for iodic acid and kM7. Figure
9 shows that this simple mechanism gave a reasonably good
fit of the experimental data for the case involving excess io-
The reaction commences with an electrophilic attack by
aqueous iodine on the electron-rich sulfur center of methio-
nine:
½R14ꢁ
R1R2S: þ IꢂI Ð ½R1R2SꢂIꢁþ þ Iꢂ
followed by the standard hydrolysis:
½R15ꢁ
½R1R2SꢂIꢁþ þ H2O ! R1R2S¼O þ 2Hþ þ Iꢂ
Addition of eqs. [R14] and [R15] gives the 1:1 experi-
mentally observed stoichiometry for this reaction. Initial
rate data, before accumulation of iodide, was used to evalu-
ate k14 = 72 1.4 (mol/L)–1 s–1.
Although the reaction is autoinhibitory, the autoinhibition
is not as strong as the one observed with the iodine–thiocya-
nate34,35 and iodine–thiourea reactions.36 In these systems, it
was assumed that iodine reacts with the product, iodide, to
produce an inert product, triiodide:
½R16ꢁ
I2 þ Iꢂ Ð Iꢂ3 ; Keq
If one assumes that the triiodide anion is inert, then the
following rate law emerges:
1
1
½1ꢁ
j
j ¼
ð1 þ Keq½IꢂꢁÞ
Rate
k14½Metꢁ0½I2ꢁ0
A plot of the inverse of the rate versus iodide concentra-
tions on the basis of eq. [1] should give a straight line, which
should allow for the calculation of k14 and a confirmation of
the well known value of Keq = 770 (mol/L)–1. This plot is
shown in Fig. 8b where the data gave a value of k14 that
was much lower than that derived from initial rate data, and
also gave a much lower Keq value that was an order of mag-
nitude lower than the literature value.37 This suggests that Iꢂ3
is not inert, and that the autoinhibition observed is derived
from the sluggish kinetics obtained with Iꢂ3 when it predom-
inates in concentrations as iodide accumulates.
Published by NRC Research Press

Chikwana et al.
697
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date concentrations, resulting in a monotonic formation of
iodine. The transient iodine formation traces (Figs. 3 and
4b) were more difficult to reproduce with this abbreviated
mechanism. It was easier to fit the initial formation of io-
dine but much more difficult to fit its consumption, even
though the model could easily reproduce the trivial iodine
consumption traces shown in Fig. 7a. The inability to cor-
rectly simulate Figs. 3 and 4b lay in our inability to cor-
rectly ascribe the correct activity of aqueous iodine in these
highly acidic solutions. Literature does not contain reliable
data on these solutions, especially when spiked with iodide.
Conclusion
Our short study has shown that iodine oxidizes the simple
biologically important thioether, methionine, to just the sulf-
oxide. Organosulfur compounds are the most effective goi-
trogenics available, and their reaction rate and mechanism
with iodine would be a significant physiological interaction.
Acknowledgements
This work was supported by grant number CHE 0614924
from the National Science Foundation.
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