Elsevier

Inorganica Chimica Acta

Volume 303, Issue 2, 30 May 2000, Pages 244-255
Inorganica Chimica Acta

The kinetics and mechanism of the oxidation of s-methyl-l-cysteine, l-cystine and l-cysteine by potassium ferrate

https://doi.org/10.1016/S0020-1693(00)00043-8Get rights and content

Abstract

The kinetics of the reactions of s-methyl-l-cysteine, l-cystine and l-cysteine with potassium ferrate were all investigated under pseudo and non pseudo first-order conditions. Methyl cysteine was oxidised to the sulfoxide and cystine was oxidised to the thiosulfonate within 300 s, with kinetics that were first-order in the concentration of the hydrogen ions, the reductant and the ferrate ions above a pH of approximately 8.4, below this pH the kinetics were independent of the hydrogen ion concentration. Under pseudo first-order conditions cysteine was oxidised within 300 ms to form the sulfinic acid with subsequent formation of a ferric–cysteine complex in excess cysteine. The kinetics involved two terms, one being first-order in the hydrogen ion, the cysteine and the ferrate ion concentrations above a pH of 8.2, and independent of the hydrogen ion concentration below this pH, and the other being first-order in only cysteine and the ferrate ion concentrations. For all three reductants the proposed mechanism involves a rate-determining step between the reductant and the protonated ferrate ion; for cysteine there is also a rate-determining step between cysteine and the unprotonated ferrate. The results are consistent with those obtained for other organosulfur and thiol reductants.

Introduction

Although potassium ferrate, K2FeO4, is a simple compound that has been known for over 100 years [1], its chemistry as an active oxidising agent has not been developed. The present study continues the investigation of the kinetics and mechanisms of the reactions between organic sulfur compounds and potassium ferrate ([2], [3], [4], [5]). Methyl cysteine and cysteine differ only in the nature of their terminal group (CH3 or H), and methyl cysteine is comparable to methionine which has one extra methylene group in its chain. Cystine comprises two cysteine moieties joined through a disulfide linkage; it is often formed as an oxidation product of cysteine.

Many of the oxidations of organosulfur compounds follow simple third-order kinetics, showing first-order dependence on the concentration of the hydrogen ions, the sulfur compound and the ferrate ions, with the kinetics becoming independent of the hydrogen ion concentration above a specific value. Some of these compounds, shown in increasing order of rate constant with the pH of the break point shown in brackets, are thioxane (8.5) [3], methionine (8.5) [2], [5] and its selenium analogue selenomethionine (8.4) [5]. Sharma and Bielski have studied the oxidation of cystine using stopped-flow and pulse radiolysis techniques at a high, constant, pH of 12.4, obtaining a second-order rate constant of 118 M−1 s−1 [6]. The suggested mechanism involves a reaction between the organosulfur compound and the protonated ferrate ion as the rate-determining step.

The thiols generally give more complex kinetics and larger rate constants. For example, 3-mercaptopropionic acid and 2-mercaptoethane sulfonic acid exhibit two kinetic terms, one of which is hydrogen ion dependent and one of which is hydrogen ion independent, whereas mercaptobenzoic acid only shows the hydrogen ion independent term [4]. In addition the decomposition of the ferrate ions is accelerated by the presence of thiols, and they form complexes with the resulting ferric ions. Preliminary studies with cysteine indicate that the reaction is too fast and complex to be studied by pseudo first-order methods in the pH range 8.9–10.2. However, Sharma and Bielski, using pseudo first-order condition at a pH of 12.4, obtained a second-order rate constant of 760 M−1 s−1 [6]. The suggested mechanism for the thiols always involves a rate-determining step between the thiol and the unprotonated ferrate ions, together with other rate-determining steps.

Methyl cysteine is oxidised first to the sulfoxide and then to the sulfone by permanganate and peroxide ions, with the latter step occurring only very incompletely under mild conditions [7]. Cystine can be oxidised by a number of reagents, particularly in large excess of the oxidant. Acidic hydrogen peroxide will cause SS fission, whereas alkaline solutions of hydrogen peroxide or permanganate will result in CS fission [8]. Cysteine is readily oxidised to cysteic acid by peroxide, or to cystine, especially in the presence of trace amounts of metals or oxygen [9]. The kinetics of the oxidation of cysteine by the chromium analogue of ferrate has been studied giving a final product of cystine and Cr(III) [10].

Section snippets

Instrumentation

The reactions were monitored spectrophotometrically with a Durrum D110 stopped-flow which was completely computer interfaced using OLIS software. Pseudo first-order kinetics were analysed with Kinfit, and non pseudo first-order kinetics were analysed using SigmaPlot 4.0. A Hewlett–Packard 8452A diode-array spectrophotometer was used to obtain reactant and product spectra. A Beckman System Gold High Performance Liquid Chromatograph, completely computer interfaced, was used for stoichiometric

The final oxidation state for iron

The problem with ascertaining the final oxidation state of iron is that, even if Fe(II) is formed as the initial product, it is readily oxidised to Fe(III), especially in the presence of even very small amounts of oxygen. Nevertheless, it is important to know the oxidation state of iron after reaction with the reductant so that the balanced stoichiometric equation can be obtained (this is especially critical for non pseudo first-order calculations). Although the presence of very small amounts

Conclusions

Potassium ferrate readily oxidises methyl cysteine to its sulfoxide and cystine to the thiosulfonate. The rate-determining step involves reaction with the protonated ferrate ion. Potassium ferrate oxidises cysteine even more readily to the sulfinate anion, with the rate-determining steps being the reaction of cysteine with both the protonated and deprotonated ferrate ions. Slight excess of cysteine forms a ferric–cysteine complex which interferes with the absorbance data. The results are

Acknowledgements

The authors thank the Goodridge Endowment, Human Resources Development (Canada), and Mount Allison University for financial support.

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