Kinetics and Mechanism of the Oxidation of N-Acetyl Homocysteine

Article pubs.acs.org/JPCA

Kinetics and Mechanism of the Oxidation of N‑Acetyl Homocysteine Thiolactone with Aqueous Iodine and Iodate Ashley Sexton,† Wilbes Mbiya,† Moshood K. Morakinyo,† and Reuben H. Simoyi*,‡ †

Department of Chemistry, Portland State University, Portland, Oregon 97207-0751, United States School of Chemistry and Physics, University of KwaZulu-Natal, Westville Campus, Durban 4014, South Africa

‡

ABSTRACT: The kinetics of N-acetyl homocysteine thiolactone (NAHT) oxidation by aqueous iodine and iodate were studied by spectrophotometric techniques. The iodate−NAHT reaction is slow and results in the formation of N-acetyl homocysteine thiolactone sulfoxide as the sole product (NAHTSO). The stoichiometry of the reaction was deduced as: IO3− + 3NAHT → I− + 3NAHTSO (S1). In excess iodate conditions, the iodide produced in S1 is oxidized to give iodine: IO3− + 5I− + 6H+ → 3I2 + 3H2O (S2). Thus in excess iodate conditions the overall stoichiometry of the reaction is a linear combination of S1 and S2 that eliminates iodide, 5S1 + S2: 2IO3−+ 5NAHT+ 2H+ → I2 + 5NAHTSO + H2O. There was a 1:1 stoichiometry for the NAHT − I2 reaction: NAHT+ I2 + H2O → NAHTSO +2I− + 2H+ (S3). All reactions, S1, S2 and S3 occur simultaneously and since they are all comparable in rate; complex dynamics were observed. Iodide catalyzes S1 and S2 but inhibits S3. Iodide is a product of both S1 and S3. It has the most profound effect on the overall global dynamics observed. The overall reaction scheme which involved S1, S2 and S3 was modeled by a simple 12-reaction mechanistic scheme which gave a very good fit to experimental data.

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chosen to work with such as thiols,16 thiocarbamides, thiocarbamates and isocyanates are the most physiologically active sulfur-based compounds. Perhaps one of the most intriguing complexities associated with sulfur chemistry is the ability of sulfur to assume a wide range of possible oxidation states. The sulfur atom, which has the ability to expand electronically into the d-orbitals, exists in a number of oxidation states ranging from −2 to +6.17 This allows sulfur to form a series of oxyanions, and other compounds at those oxidation states that are chemically and biologically active. Such compounds can undergo redox reactions with energy being liberated as the oxidation state of sulfur increases. Electron flow from an electron donor via several intermediate electron carriers toward an electron acceptor is the means by which biochemical energy is generated.11 These carriers are located in an orderly manner at and within membranes, and the electron transport across these membranes gives rise to potential differences, which in turn liberate energy for metabolic processes. Sulfur compounds in higher oxidation states act as sinks for such reactions. Reduced sulfur compounds are excellent electron donors, while the intermediates between sulfide and sulfate can play both

INTRODUCTION Thiols have been studied extensively for their free radical scavenging and antioxidant properties that neutralize reactive oxygen species responsible for damage to proteins, carbohydrates, lipids, and DNA.1,2 They are found in a wide range of biological molecules including the amino acids cysteine and glutathione, as well as in some polysaccharides, and lipids.3 −2, 0, and +6 are the most stable oxidation states for sulfur centers, although it is possible to achieve other states in between. The roles of most inorganic metals in biological chemistry can be easily studied and rationalized. This includes metals such as iron, molybdenum and chromium.4−10 It has, however, not yet been possible to derive a realistic role for sulfur in physiological mechanisms. Sulfur is the seventh most abundant element in the known universe, and the eighth most abundant constituent of living organisms.11,12 It is essential for the growth and proliferation of all living organisms from microbes to humans, and forms an integral part of living structures.13 The stochastic nature of sulfur chemistry due to prolific free radicals formation has precluded the elucidation of a generic oxidation pathway of a sulfur center to the extent that oxidation mechanisms could vastly differ on the basis of very slight substitution differences as well as on reaction conditions14,15 By working with small organosulfur compounds, we have been able to control sulfur− sulfur polymerizations to at most the dimeric species only. Fortuitously, these small organosulfur compounds we have © 2013 American Chemical Society

Received: August 26, 2013 Revised: October 27, 2013 Published: October 28, 2013 12693

dx.doi.org/10.1021/jp408540u | J. Phys. Chem. A 2013, 117, 12693−12702

The Journal of Physical Chemistry A

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roles depending on the situation.11 There are a vast number of compounds in which one or more oxygen atoms are linked to the sulfur center. These include sulfoxides, sulfones, sulfamates, as well as the well-known sulfur oxo-acids (sulfenic, sulfinic, and sulfonic acids); all of them physiologically active. The chemical and biological properties of organic sulfur−oxygen compounds depend on the nature of the substituents on the sulfur, carbon, and oxygen atoms. The character of the C−S bond will be strongly influenced by the substituents on both the carbon and the sulfur atoms.18 Furthermore, the medium (polarity, ionic strength, temperature and pH) will also exert a very powerful effect. Sulfur centers are especially susceptible to oxidation due to their nucleophilic properties resulting in a series of oxo-sulfur acids: sulfenic, sulfininc, and sulfonic acids.19 Oxyhalogen organic sulfur chemistry has often demonstrated nonlinear kinetics that can be explained by invoking autoinhibition, autocatalysis, autoxidations, clock reactions behavior, oscillations, and transient halogen formation. Of the halogens, iodine has been extensively studied for its dependence in proper thyroid function. A slew of health problems can be caused by iodine deficiency, most commonly, hypothyroidism, but alarmingly it is the leading cause for preventable retardation in infants and unborn fetuses.20,21 Present throughout the body, iodine is found among living cells and is often oxidized in biological systems.

hopes to broaden the understanding of potential future pharmaceutical uses.

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METHODS The following reagents were used without further purification: all D-L-N-acetyl homocysteine thiolactone (99%, Acros), sodium iodide, sodium perchlorate, sodium iodate, iodine (Sigma-Aldrich), perchloric acid, 70%, (Fisher Scientific). Saturated iodine solutions were prepared by dissolving iodine crystals in distilled deionized water with stirring for at least 12 h. Excess iodine was removed by vacuum filtration. The solution was protected from light by keeping the flask capped and covered with aluminum foil and stored in the dark. Iodine concentrations were determined spectrophotometrically at 460 nm using 770 M−1 cm−1 for the absorptivity coefficient. Ionic strength for all reaction was maintained at 1.0 M (NaClO4). For the iodate−NAHT system, the stoichiometry was determined spectrophotometrically by observing the ratio of iodate to NAHT where iodine concentration began to be completely consumed after 48 h. The experimental determination was performed as follows: A number of stoichiometric determination experiments were carried out at constant NAHT concentrations while varying iodate concentrations. After 48 h, excess oxidizing power is determined iodometrically and the final iodine absorbance spectrophotometrically determined. A plot was then made of thiosulfate titer or iodine absorbance versus initial iodate concentrations. These plots are extrapolated to zero titer or zero absorbance. The value of the intercept on the mantissa represents the iodate concentrations needed for complete consumption of NAHT with no excess iodate left to effect the Dushman reaction and produce iodine. This will be the stoichiometry of the reaction. For the iodine−NAHT reaction, two methods were used to determine the reaction stoichiometry. In the first method, varying amounts of excess iodine over NAHT were used. After 6 h, the residual absorbance at 460 nm was determined. These absorbances were next plotted against initial iodine concentrations, with the intercept on the iodine axis representing the stoichiometry of the reaction. The other method involved a titration of iodine versus NAHT. Iodine was in the buret. Indicator was freshly prepared starch. Stoichiometry was determined by the point where the blue-black color lingered.

N-acetyl homocysteine thiolactone (NAHT), commonly known as citiolone, is the cyclic form of homocysteine with a blocked α-amino group. Known as an antioxidant and free radical scavenger, it does not exist naturally in the body but has potential to have a multitude of pharmaceutical uses including the treatment of hepatitis.22 It also has been shown to increase intercellular glutathione levels in skin exposed to UV radiation, and has prevented glutathione deficiencies normally caused by thallium and phosphamidon toxicity.23−25 Treatment with NAHT has been observed to effect an increase in pancreatic islets cells normally destroyed by diabetes and has prevented growth malformations in diabetic rat embryos.26,26 Interesting enough, neuronal lipofusion, the “aging pigment,” has been shown to be inhibited by NAHT, and it has been shown to reduce the appearance of wrinkles by thickening the skin when applied topically.27 It has commonly been used in the thiolation of proteins by reacting with primary amines in a nucleophilic ring-opening reaction without damaging the structure of the protein and is often followed by acylation.28−32 Influenced by this ring-opening reaction, NAHT was successfully used in an amine-thiol-ene conjugation for a simple, one-pot photopolymerization process.33 Plausible mechanisms of methionine, cysteamine, and aminoiminomethanesulfinic acid with iodate or aqueous iodine have been determined in previous studies, although none with NAHT.34−39 In this study, the kinetics and mechanism of NAHT with aqueous iodine and iodate were investigated in

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RESULTS AND DISCUSSION Stoichiometry. All our stoichiometric determinations pointed to the production of a single product that involved the addition of a single oxygen atom to NAHT coupled with a 2-electron oxidation of NAHT. No other products were observed. Figure 1 shows the titrimetric analysis of the acidic iodate−NAHT reaction in excess iodate. The intercept value on the iodate concentration axis suggests a 1:3 stoichiometry. At this 1:3 stoichiometry, no iodine is formed, and the final reduction product of iodate is iodide. The iodine center undergoes a 6-electron reduction, meaning each NAHT molecule loses 2 electrons to form a sulfoxide or an S-oxide. ESI spectra (vide infra, Figure 2) clearly indicate that this product is a sulfoxide in which the oxygen atom attaches to the sulfur atom. In this manuscript, we will assume this sulfoxide is NAHTSO: IO3− + 3NAHT → 3NAHTSO + I−

(R1)

The reaction scheme is shown below as Scheme 1. 12694



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presence of strongly electron-donating groups. The thiolactone molecule is neither of these. The ESI-MS spectrum shown in Figure 2 was obtained in the negative mode. It was obtained in stoichiometric excess of NAHT such that at the end of the reaction we would expect to obtain both the substrate NAHT and its oxidation product, NAHTSO. Spectrum only shows two peaks, one at m/z = 158.03 belonging to NAHT and the other at m/z = 174.89 belonging to the sulfoxide. NMR spectra were also obtained of NAHT and NAHTSO. They showed a shift in the multiplex peak from the proton at the asymmetric from 4.46 to 4.60 ppm. The protons next to the oxidized sulfur center shifted from 2.12 and 2.54 ppm to 2.24 and 2.69 ppm, respectively, due to the addition of the electron-withdrawing oxygen atom on the sulfur center. The iodine−NAHT reaction gave a stoichiometry of 1:1, also indicating a two-electron oxidation of NAHT. The ESI-MS spectrum for the product of the reaction between iodine and NAHT shows only two peaks, one at m/z = 174.89 for the product, and another at m/z = 126.98 for iodide. In excess iodate conditions, the iodide produced in stoichiometric reaction R1 will react with the excess iodate to produce aqueous iodine as expected through the Dushman reaction:

Figure 1. Iodometric titration of excess oxidizing power after complete consumption of NAHT in excess iodate conditions. NAHT concentrations were fixed at 6 mM. The intercept value is 1.808 mM iodate for a complete consumption of NAHT with no iodate left to produce iodine. This suggests a 1:3 stoichiometry.

Formation of a sulfenic acid is highly unlikely. It would entail the formation of a CS double bond, which would disrupt the stable five-membered ring of the thiolactone; while with the sulfoxide, the sulfur center still retains a tetrahedral symmetry due to the presence of a lone pair of electrons on the sulfur center. Sulfenic acids have traditionally been stabilized only in conditions of highly stearically hindered sulfur centers or in the

IO3− + 5I− + 6H+ → 3I 2 + 3H 2O

(R3)

Thus the stoichiometry of the reaction in excess iodate is a linear combination of reactions R1 and R3 that consumes all

Figure 2. The mass spec analysis run in the negative mode where relative abundance is plotted against mass to charge ratio. The base peak of 174.89 represented the NAHTSO, and the second largest peak at 158.03 is the NAHT. 12695



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Scheme 1. Reaction R2

commenced rapidly and slowed down until the final expected iodine formation was attained. This is in contrast to the slow initial rates observed in Figure 3a,b. The variation of NAHT displayed first order rate dependence behavior with an initial rapid iodine formation followed by a constant rate. No change in the induction time was observed. In these traces, NAHT does act as the limiting reagent in iodine production, where, in theory, the final absrobance would refect a concentration of 0.2 times the initial NAHT concentration. The first-order rate dependence confirmed the linear regression observed in plotting initial rates against NAHT concentration. Iodine NAHT Dynamics. The direct reaction of iodine with NAHT is the most important reaction in the overall reaction scheme. It provides both positive and negative feedback in the iodate−NAHT reaction. Its rate of reaction will determine whether a monotonic increase in iodine formation is observed, or a transient formation. If it is orders of magnitude faster than the iodate−NAHT reaction, then it should be able to mop up all iodine formed via the Dushman reaction. Thus observation of iodine would indicate that all the NAHT has been totally consumed. As a positive feedback, its product, iodide, can feed back into the oxidation of NAHT by promoting formation of reactive species HIO2 and HOI from iodate (vide infra). Figure 4a shows that the rate of this reaction is comparable with the fiduciary iodate−NAHT reaction. Figure 4a is an iodine concentration dependence series of experiments. Starting absorbances, as expected, are determined by the initial iodine concentrations used. Figure 4b shows that this reaction is first order in iodine. The plot of initial rate versus iodine concentrations gave a linear relationship with an intercept kinetically indistinguishable from zero, suggesting that that there is no other pathway for the consumption of iodine except for its direct reaction with NAHT. Figure 5a shows a series of experiments performed at constant iodine concentrations while varying [NAHT]0. NAHT was in excess for the whole series of experiments. At reaction completion, expected iodine concentrations will be zero. The general appearance of the absorbance traces in Figure 5a is one of an initial relatively rapid reaction, followed by a general slowing down as the reaction proceeds. While this is a hallmark of second order kinetics, there also appears to be some degree of autoinhibition. Figure 5b shows the reaction as being first order in [NAHT}0. Experimental data from Figures 4 and 5 allowed for the derivation of a bimolecular rate constant of the direct reaction between iodine and NAHT of k = 1.16 ± 0.21 M−1 s−1. The effect of acid was difficult to determine and assess. Figure 6 shows that acid has a small inhibitory effect on the reaction. As acid is increased, the equilibrium of the following hydrolysis reaction is pushed to the left:40

the iodide produced in R1 and convert it to aqueous iodine. This linear combination is 5R1 + R3: 2IO3− + 5NAHT + 2H+ → I 2 + 5NAHTSO + H 2O (R4)

The reaction was studied exclusively in this environment since formation of iodine was utilized to follow the reaction. At these conditions, final iodine produced was determined by initial NAHT concentrations, and was exactly 20% of the initial NAHT concentrations. If we assume an oxidant to reductant ratio, R as [IO3−]0/[NAHT]0; then stoichiometry R1 is obtained when R ≤ 0.33 and stoichiometry R4 is obtained when R ≥ 0.40. Mixed stoichiometry is obtained when 0.33 < R < 0.40. Reaction Kinetics. In general, the reaction is very slow, often taking 8 h to proceed to completion. Only those reactions undertaken to decipher the stoichiometry were carried out to completion in which reaction solutions were incubated for 48 h or more. Most reactions were carried out only to as far as 2500 s. Contrary to most oxyhalogen−sulfur reactions, the iodate− NAHT reaction displays no induction period. Iodine formation is observed as soon as reagents are mixed together (see Figure 3a−c). The variation of iodate concentration with constant acid and NAHT showed sluggish reaction dynamics. The rate of iodine production increased as the iodate concentration increased. The reaction was initially very slow, but its rate seemed to increase after the absorbance exceeded 0.02. This is a competitive mechanism with iodine observation being a direct result of the reactions that produce iodine exceeding those reactions that consume it. If the concentration ratios of NAHT to iodate were conducted at stoichiometric ratios or if NAHT was in slight excess, only transient iodine formation would be observed. By plotting initial rates against iodate concentration, the reaction was determined to be first order dependent in iodate. This result was only qualitative. Instantaneous rates of reactions were taken 20 s into the reaction reactions for a qualitative determination of the effect of iodate. Acid is not a reactant in the stoichiometry of the reaction, but acts as a powerful catalyst (Figure 3b). All traces shown in Figure 3b gave the same final iodine absorbance. Using the same qualitative approach utilized for iodate dependence, a square acid dependence on the rate of reaction was observed. All experimental traces shown in Figure 3c involved excess of iodate, with R > 0.40. As expected, final iodine concentrations observed were determined by the initial NAHT concentrations. Rate of formation of iodine was dependent on NAHT concentrations to the first power. At these high oxidant-toreductant ratios and highly acidic conditions, iodine formation 12696



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Figure 4. (a) Traces from varied iodine concentrations show an increase of rate as the initial iodine concentration increases. Initial NAHT concentration was held constant at 0.001M, where [I2]o was (a) 1.9 × 10−4, (b) 2.3 × 10−4, (c) 2.8 × 10−4, (d) 3.7 × 10−4, (e) 4.7 × 10−4 M. (b) The rate dependence on iodine was determined to the first power as suggested by the linear regression deduced by plotting rate against concentration.

I 2 + H 2O → HOI + I− + H+; KH

(R5)

Hence the initial iodine concentrations in Figure 6 increase with increase in acid concentrations. A plot of inverse of reaction rate with [H+]0 was linear (plot not shown). The effect of iodide on the reaction was also determined. Although iodide was strongly inhibitory, there was no simple relationship that could be derived. This can result from a reaction with a multiterm rate law. Mechanism. Data derived from Figure 3a−c imply the strong dominance of the standard oxyiodine kinetics as being dominant in the initiation and subsequent propagation of the reaction. Iodate, itself, is very inert. The initial steps in any acidic oxidation by iodate is the formation of the reactive oxidizing species, HIO2 and HOI:37,38,41

Figure 3. (a) Monitoring the iodine production from the effect of iodate on the oxidation of NAHT . [NAHT]o = 0.0015 M, [H+]o = 0.1 M, [IO3−] = (a) 0.01, (b) 0.02, (c) 0.03, (d) 0.04, (e) 0.05, (f) 0.06, (g) 0.07 M. (b) Absorbance traces for the variation of acid. [NAHT]o and [IO3−]o were 0.0015 and 0.05 M, respectively, while acid was varied from (a) 0.050, (b) 0.075, (c) 0.085, (d) 0.100 M. Acid acts as a strong catalyst, and also influences induction time. (c) Absorbance traces of the oxidation of NAHT by observing iodine. The iodate and acid were both in excess with initial concentrations of [IO3−]o = 0.1 M [H+]o = 0.2 M while [NAHT] varied from (a) 0.001 M to (f) 0.006 M in increments of 0.001 M.

IO3− + 2H+ + 2e− ⇄ HIO2 + OH−

(R6)

Reaction R6 has been extensively studied and, if the two electrons are sourced from iodide, it is the initiation reaction for the Dushman reaction:39,42 IO3− + 2H+ + I− ⇄ HIO2 + HOI 12697

(R7)



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In reaction R7, iodide is oxidized from oxidation state of −1 to +1 in HOI, while the iodine center in IO3− is reduced from +5 to +3 in HIO2. Reaction R7 is a composite of at least four steps which involve an attack by iodide (or any appropriate nucleophile) on a protonated iodate species: HIO3 + I− ⇄ HI 2O3−

(R8)

The rate-determining step is an acid-assisted break-up of the intermediate HI2O3−: HI 2O3− + H+ ⇄ HIO2 + HOI

(R9)

With the assumption of R9 as rate-determining, then the initial rate of reaction is Rate = k 0[IO3−]0 [I−]0 [H+]0

(1)

Figure 3a−c suggests the following rate of reaction: Rate = k 0[IO3−]0 [NAHT]0 [H+]0

(2)

Thus it is plausible to assume that the two electrons, as depicted in reaction R6, are derived from NAHT, since in general, iodide ions are not initially added to the reaction mixture, and have to be generated as the reaction proceeds: HIO3 + NAHT + H+ → HIO2 + NAHTSO + H+ (R10) +

Reaction R10 is correct as written: H is absolutely essential on the left-hand side of reaction R10; the expulsion of a proton on the right-hand side is due to the preferred formation of the sulfoxide over a sulfenic acid, which would be unstable.43 There would be a cascade of two further reactions from R10 in which iodous acid is finally reduced to iodide: Figure 5. (a) [NAHT] was varied from (a) 0.004 to (e) 0.0012 M in increments of 0.002 M, while [I2]o was constant at 4.15 × 10−4 M. As the initial concentration of NAHT increased, so did the rate of oxidation. (b) The dependence on NAHT was determined to be first order. Intercept was also kinetically indistinguishable from zero for this plot.

HIO2 + NAHT → HOI + NAHTSO

(R11)

HOI + NAHT → H+ + I− + NAHTSO

(R12)

The iodide produced in R12 can be recycled in R7 to fuel the initiation reaction of the Dushman reaction. Addition of reactions R7, R11, and R12 and applying the appropriate stoichiometric factors will show that two iodide anions will eventually produce three iodide anions as in cubic autocatalysis kinetics. Thus we expect a continuous increase in iodide ions as the reaction progresses, even in reaction environments in which iodide anions had not been deliberately added to the reaction mixture. This will make reaction R7 the overall ratedetermining step. Our experimental data confirm this fact. Iodide, as has been observed in our experimental data, catalyzes the overall reaction by removing the need for the R10−R12 cascade, and immediately making R7 relevant. Iodine−NAHT Reaction. The kinetics of this reaction were more complex than the standard bimolecular kinetics suggested: Rate = −

d[I 2] = k 3[NAHT][I 2] dt

(3)

The initial part of the iodine−NAHT reaction is an electrophilic attack by iodine on the nucleophilic sulfur atom of the thiolactone for a two-electron oxidation of the sulfur atom from oxidation state −2 in the thiolactone ring, to 0 in the sulfoxide:

Figure 6. Acid variation with the oxidation of NAHT with iodine. Acid was shown to inhibit the reaction over a range of (a) 0.122, (b) 0.161, (c) 0.200, (d) 0.239, and (e) 0.278 M, while initial NAHT and I2 concentration were held constant at 0.01 and 5.1 × 10−4 M, respectively.

NAHT + I 2(aq) → NAHTSO + 2H+ + 2I− 12698

(R13)



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The rate-determining step is the initial electrophilic attack by iodine; the subsequent hydrolysis should be rapid and facile. This hydrolysis is such a dissociative process that entropy terms for the reverse reaction are prohibitive and reaction is essentially irreversible. Further analysis of this reaction needs to address the inhibitory effects of iodide (a product) and acid. As reaction R13 proceeds, the iodide generated reacts rapidly with remaining aqueous iodine to produce triiodide, I3−:44−46 I 2 + I− ⇌ I3−;

Keq = 770 M−1

( (

[I 2]0 1 + Keq[I−]

(R14)

⎛ k K ⎞ −d[I 2] = [NAHT][I 2]o ⎜k13 + −12 H+ ⎟ ⎝ dt [I ][H ] ⎠

Values of [I ]t are dynamic and vary with time. k13 was determined in this study, and k12 was determined from the best simulations fit (vide infra). Figure 6 only displays very mild acid retardation over a 0.15 M change in acid concentrations. The forward rate constant of reaction R5 as it is written is 2.2 s−1, and its reverse is 3.1 × 1012 M−2 s−1.40 Thus it was easy to evaluate the relative changes in the concentrations of HOI and I2 over the range of acid concentrations used for Figure 6. Between 0.122 M H+ and 0.278 M, there was effectively no change in the ratio of [HOI] to [I2]. Thus acid retardation could not be ascribed to the perturbation of the equilibrium of reaction R5. To further support this fact, within the range of acid concentrations utilized in our study, coupled with the fact that KH is so small; the second term in the brackets of eq 10 is never at any time significant. The other possible source of acid retardation could be in the protonation of the sulfur center of the thiolactone, thus making it less nucleophilic. Reaction R13 relies on the nucleophilicity of the sulfur center to enhance its rate. We could mathematically derive a rate of reaction that can include acid retardation using the same argument utilized for iodide retardation. We assume, again, in the extreme case, that the protonated thiolactone is not reactive, and that only the unprotonated form would undergo reaction R13. Let us ascribe to the thiolactone, a basicity parameter, KB;

(4)

(5)

(6)

In the limit of low iodide concentrations, eq 6 reverts to familiar eq 3. Equation 6 can now justify the observed iodide inhibition. Equation 6 is derived from the assumption that I3− is inert, and it predicts, in the limit of high iodide concentrations, that the reaction would shut down completely. This extreme inhibition by iodide was not experimentally observed. I3− is thus reactive, but not as actively as aqueous iodine. Acid retardation could arise from two possible sources. One is related to reaction R5 and has been mentioned previously. If there is a marked difference in the reactivities of I2 and HOI, then pH variations would impact on the rate. If rate of oxidation by HOI is more facile than by I2, then an increase in acid would push equilibrium of reaction R5 to the left, depleting [HOI] while increasing [I2] and thus lowering the rate. Let us assume two pathways, reactions R12 and R13 with Rate = −

NAHT + H+ ⇄ [NAHTS−H]+ ; KB

[NAHT] =

[NAHT]0 1 + KB[H+]

(11)

where [NAHT]0 is the initial NAHT concentration that has now been distributed into unprotonated and protonated forms. If we combine both iodide and acid retardation, we obtain a final rate eq 12: k 3[NAHT]o [I 2]o −d[I 2] = dt (1 + KB[H+])(1 + Keq[I−])

(7)

We can use the hydrolysis equilibrium constant of reaction R5 to eliminate [HOI] in eq 7 and obtain eq 8: ⎛ k K ⎞ −d[I 2] = [NAHT][I 2]⎜k13 + −12 H+ ⎟ ⎝ dt [I ][H ] ⎠

(R15)

Utilizing the same mass balance technique,

d[NAHT] dt

= k12[HOI][NAHT] + k13[I 2][NAHT]

(10)

−

Rate eq 3 can now be rewritten as k [NAHT][I 2]o −d[I 2] = 3 dt 1 + Keq[I−]

(9)

We know, from literature values, that KH = 4.13 × 10 . Our prevailing iodide concentrations were in the ranges of 10−4 M. These concentrations make the bottom term insignificant because 1 ≫ KH/[I−][H+] making the bottom term unity, unless the pH is above 8 (we worked around pH 1 and 2). Equation 9 then reduces to

Leading to [I 2] =

) ) −12

Thus from reaction R14, iodide inhibition can be rationalized by the conversion of iodine to triiodide, a less-effective electrophile. The new rate law can be derived as follows. A mass balance equation exists that links the iodine species, coupled with the R14 equilibrium. Aqueous iodine concentrations can then be expressed in terms of total iodine concentrations, [I2]0 and iodide. Instantaneous iodine concentrations at any time during the reaction, [I2], can thus be expressed as [I 2] = [I 2]0 − Keq[I 2][I−]

k K

k13 + [I−12][HH+] −d[I 2] = [NAHT][I 2]o KH dt 1 + [I−][H + ]

(12)

Equation 12 is only approximate. KB is not known and is best derived from the best fit to the kinetics data. One would expect this value to vary with pH and even with ionic strength of the medium. In the limit of high pH and low iodide conditions, eq 8 is approximated by eq 3. Knowing that high pH environments do not contain aqueous iodine, it is then important to note that eq 8 applies within a small range of acid concentrations within which a value of KB will be derived from observed kinetics data.

(8)

We can then institute a mass balance equation on the iodine species so that the rate is with respect to the initial concentrations of iodine, [I2]0: 12699



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Table 1. Proposed Reduced Overall Reaction Mechanism for the Iodate NAHT Reaction number

kf; kr

reaction

M1

IO3− + I− + 2H+ ⇌ HIO2 + HOI

2.8; 1.44 × 103

M2

HIO2 + I− + H+ ⇌ 2HOI

2.1 × 108; 90

M3

HOI + I− + H+ ⇌ I 2 + H 2O

3.1 × 1012; 2.2

M4

IO3− + HOI + H+ ⇌ 2HIO2

8.6 × 102; 2.00

M5

I 2 + I− ⇌ I3−

6.2 × 109; 8.5 × 106

M6

IO3− + H+ ⇌ HIO3

2.04 × 108; 1.25 × 109

M7

HIO3 + NAHT + H+ → HIO2 + NAHTSO + H+

8.8 × 10−2

M8

IO3− + NAHT + H+ → HIO2 + NAHTSO

1.05 × 10−4

M9

HIO2 + NAHT → HOI + NAHTSO

50

M10

−

HOI + NAHT → NAHTSO + I + H

+

125

+

M11

I 2 + NAHT + H 2O → NAHTSO + 2H + 2I

M12

I3−

+

−

+ NAHT + H 2O → NAHTSO + 2H + 3I

−

1.158 0.785

Legend: Forward and reverse rate constants separated by a semicolon. Units determined by reaction molecularity except where water is involved.

Overall Reaction Scheme. The overall reaction scheme combines three reactions which are relevant during the reaction’s duration: (a) the iodate−NAHT reaction; (b) the iodate−iodide reaction, and (c) the iodine−NAHT reaction. Since the reaction does not show any induction period, it shows that reaction (c) is too sluggish to mop up all the iodine formed by reaction (b). Iodine is formed in only one reaction, the reverse of its hydrolysis, R5. The overall rate of formation of iodine is the difference between the reverse of its hydrolysis and reaction R13 d[I 2] = v−5 − v13 > 0 dt

formation and consumption of iodine. The central position of iodide was evident in our attempts at modeling the reaction. The simulations were undertaken as if no iodide was initially added to the reaction mixture. We initially assumed that initiation of the reaction was through the trace amounts of iodide of about 10−6 M that exist in all iodate solutions. Iodide was then allowed to build up through the in-built cubic autocatalytic production. This could be achieved by setting forward rate constant of M7 to zero. Without iodide generated from the M7−M10 cascade, the model generated much slower kinetics than those experimentally obtained. Thus, the bulk of the iodide ions needed for reaction (a) to proceed were derived from the M7−M10 cascade. Not surprisingly, the most important parameter for the whole reaction scheme was kM7, the forward rate constant of the reaction of acidified iodate with NAHT. It, effectively, was the only adjustable parameter since kinetics parameters for M1−M6 are effectively literature values and M11 (and, implicitly, M12) were derived from this study. Slight alterations to kM7 resulted in marked changes in the model profile. The emergence of kM7 as the effective ratedetermining step in the model retained the kinetics rate law as in eq 1. The next important kinetics parameter was kM11. Though this was fixed, having been experimentally determined as 1.158 M−1 s−1; it was important that it was retained at this low value to allow for the immediate accumulation of iodine at the start of the reaction with no discernible induction period. kM12 was maintained lower than kM11 based on the observed iodide inhibition. What was remarkable was how unimportant kM10 was to overall reaction dynamics. The global reaction dynamics only changed slightly with orders of magnitude changes in the value used for the model. It was important in the generation of iodide for reaction propagation only. Two reasons for its lack of importance: (a) in highly acidic conditions, such as those employed for this study, HOI concentrations never raised to above 10−8 M and (b), in high acid, reaction R5’s equilibrium lies overwhelmingly to the left, and iodine is the preferred iodine species over HOI. Figure 7 shows how well this simple model mimicked the experimental traces. The model gave a very reasonable fit and was able to predict both acid and iodate dependence experimental data.

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The overall reaction scheme is summarized in Table 1, and has been reduced to a network of only 12 reactions. The first six equations have been previously studied with well-known kinetics parameters: oxyiodine reactions (M1−M5), and the rapid I2/I− equilibrium (M6). Reactions M7−M12 are the possible reactions that involve NAHT with an oxyiodine species. All these reactions were assumed as being essentially irreversible. Reaction M6, the association/dissociation of iodic acid, was derived from the silver iodate solubility studies of Naidich and Ricci.47 Although a dissociation constant for HIO3 of 0.163 was derived, it was not easy to extrapolate to our conditions (pH and ionic strength). This is a rapid, diffusioncontrolled protolytic process. The values adopted for the forward and reverse sections of the reactions were high, with the only crucial prerequisite being adherence to the dissociation constant of iodic acid. The overall reaction network was greatly simplified by the presence of only one product, NAHTSO. Thus, only one reductant existed in the reaction solution, with five possible oxidants: IO3−, HIO2, HOI, I2 and I3−. Iodide is the most crucial intermediate and/or product in the whole reaction mechanism. It is involved in all three reactions that make up the overall scheme. It is catalytic in reaction (a) and inhibitory in (c). It, however, does not act as a classic autocatalyst or autoinhibitor. Its effect on (a) is linear and on the basis of linear mass-action kinetics. Computer Simulations. The mechanism was modeled using Kintecus modeling software designed by James Ianni.48 This simple network of 12 reactions incorporated both the 12700



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Figure 7. Simulations of iodate variations for 2000 s using Kintecus. Traces a and b correlate with traces d and g from Figure 3a, respectively.

Traces d and g were simulated from Figure 3a for 2000 s. The model was less accurate over longer periods of time, most likely due to the activity of aqueous iodine, which is unknown at our reaction conditions. Aqueous iodine itself saturates at approximately 0.9 mM, with expected maximum absorbances in the regions of circa 0.700. Iodine can rise above 0.9 mM in the presence of iodide, which produces highly soluble triiodide, I3−

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AUTHOR INFORMATION

Notes

The authors declare no competing financial interest.

ACKNOWLEDGMENTS This work was supported by Research Grant Number CHE 1056311 from the National Science Foundation. REFERENCES

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