Peroxodisulphate ions, reactions

The rate equation derived from this mechanism is in accord with most of the observed features, but it predicts that with excess substrate the second-order rate coefficient should decrease during a run, whereas the observed rate coefficient always increases during a run, irrespective of whichever reactant is in excess. Whalley et al. suggest that incomplete dissociation of peroxodisulphate in the solvent might be responsible for the discrepancy. Another discrepancy is pointed out by Wilmarth and Haim, but these authors agree with Whalley et al. in concluding that the initiation step is reaction (91) rather than the spontaneous fission of the peroxodisulphate ion. 

Studies on the kinetics and mechanism of the oxidation of ascorbic acid show the rate to be independent of the concentration of organic substrate. A chain mechanism is proposed with reaction between the radical anion of ascorbic acid and the peroxodisulphate ion. Although the influence of neutral salts is negligible and that of acid slight, the presence of allyl acetate strongly inhibits the rate, suggesting the formation of radical ions. In the reaction with lactic acid, a similar rate expression is observed although the description of the mechanism is different. 

Finally, there is an extremely violent reaction during the action of ammonia on ammonium peroxodisulphate in the presence of Ag- ions. 

One drop of dilute silver nitrate has to be added to speed up the reaction. Silver ions act as catalysts the catalytic action is due to the transitional formation of silver(III), Ag3+. Halides must be absent they can be removed easily by evaporating the solution with concentrated sulphuric acid until fumes of sulphur trioxide appear. After cooling, the solution can be diluted and the test carried out. The excess of peroxodisulphate can be decomposed by boiling ... 

In the acid-catalysed process, dominant for pH less than 2, the products isolated may be oxygen (below 0.5 Af acid), peroxomonosulphuric acid, a mixture of the two, or may contain detectable amounts of hydrogen peroxide or ozone. The formation of hydrogen peroxide (at 98 °C in 2 M HjSO ) has been examined and shown to be catalysed by a number of metal ions. Under all conditions examined the oxygen liberated arises entirely from the peroxide . Exchange between S-labelled sulphate and peroxodisulphate proceeds very slowly, if at all, in this region . Further, it appears probable that this particular decomposition path does not involve intermediates which react at appreciable rate with cerium-(III) It has been suggested that the initial, overall, reaction is... 

Peroxodisulphuric acid, H2S2O8, is a strong acid whose second pK is below zero (Kolthoff and Miller ). Under the conditions normally employed in peroxodisulphate oxidations (aqueous solution, pH >1) the ion 8208 is the dominant species. The ion is a powerful two-electron oxidising agent with a redox potential of — 2.01 V. In the majority of its reactions the primary step is the formation of sulphate radical-ions, either by spontaneous fission of the peroxide bond, or by attack on a substrate X, i.e. 

Many of the reactions of peroxodisulphate are catalysed by silver ions, and although the reactions then involve higher oxidation states of silver as oxidants, it is convenient to consider them along with the uncatalysed oxidations. Silver ions also catalyse the decomposition of peroxodisulphate in a second-order reaction, viz. 

The possible types of chain mechanisms for peroxodisulphate oxidation have been classified by Wilmarth and Haim according to the dominant initiation and termination steps, and the relative importance of sulphate radical-ions and hydroxyl radicals in the propagation steps. Some of the rate equations corresponding to the different types of mechanisms are the same, so the observation of a particular rate equation does not always permit a unique mechanism to be inferred. In certain cases the nature of the chain initiation step can be deduced from the effect of a free-radical scavenger on the reaction rate. Thus in the oxidation of 2-propanol, the addition of allyl acetate reduces the rate to that observed for the spontaneous decomposition of peroxodisulphate, indicating that the chain initiation step is the same as the rate-determining step of the spontaneous decomposition, viz. the fission of peroxodisulphate into sulphate radical-ions. 

Later work showed this mechanism to be incorrect. Wiberg showed that 804 when present in the reaction mixture does not give labelled peroxodi-sulphate, as required by the reversible first step of Levitt and Malinowski s mechanism. Furthermore, allyl acetate inhibits the reaction and reduces the rate of consumption of peroxodisulphate to that observed in the absence of 2-propanol. Wiberg proposed a chain mechanism involving sulphate and hydroxyl radicals. In a thorough study of the reaction, Ball et aV showed that all previous studies were complicated by the catalytic effects of trace amounts of metal ions (most likely cupric ions) and inhibition by dissolved oxygen from the atmosphere. In the absence of oxygen there is no catalysis by cupric ions, and the rate equation is... 

Dogliotti and Hayon generated sulphate radical-ions at room temperature by the flash photolysis of peroxodisulphate solutions, and measured the rate at which they attack 2-propanol, i.e. the rate of reaction (21). They found kj = 8.5 3.0xl0 l.mole.sec" at pH 4.4. 

The reaction is faster than the spontaneous decomposition of peroxodisulphate, so Bartlett and Cotman proposed a chain mechanism involving sulphate radical-ions and the CH2OH radical. Kolthoff et al. showed that allyl acetate inhibits the reaction, and reduces the rate to that observed in the absence of methanol. They pointed out that if the inhibition is explained on the basis of Bartlett and Cotman s mechanism, the predicted rate equation does not include the methanol concentration. This difficulty was resolved by Edwards et al., who showed that in the absence of oxygen the reaction is zero-order with respect to methanol. They proposed the following mechanism (similar to that originally proposed by Bartlett and Cotman)... 

In the absence of silver ions the oxidation is extremely slow. The reaction is first-order with respect to both peroxodisulphate and silver ions, and the rate is independent of the pinacol concentration. Menghani and Bakore suggest a chain mechanism involving sulphate radical-ions and silver(II) ions. The same workers found the same type of rate equation for the oxidations of propane-1,3-diol and butane-l,3-diol, and propose the same mechanism. 

Subbaraman and Santappa studied the oxidations of formaldehyde and acetaldehyde in de-aerated solutions, both in the presence and absence of silver ions. When the concentration of aldehyde is much less than that of peroxodisulphate, the rate equation for reaction in absence of silver ions is... 

Srivastava and Ghosh report that the kinetics are first-order with respect to peroxodisulphate and zero-order with respect to formic acid, but Kappana reports first-order kinetics with respect to each reactant. The effect of trace amounts of metal ions and of oxygen on the rate is uncertain, and discussion of the mechanism is of doubtful significance at present. However, the reported observations definitely indicate a chain mechanism. Thus Srivastava and Ghosh found an induction period in the oxidation, and report that halide ions inhibit the reaction (inhibition by halide ions is a feature of reactions involving hydroxyl radicals). In a study of the silver ion-catalysed oxidation, Gupta and Nigam found that the reaction is approximately first-order with respect to both peroxodisulphate and the catalyst, and zero-order with respect to the substrate. 

The reaction is very sensitive to metal ion catalysis, particularly by Cu " and Ag", and oxygen inhibits the reaction. Po and Allen studied the uncatalysed reaction in oxygen-free solutions containing 10 M EDTA to ensure that the concentrations of free metal ions were insignificant. Under these conditions the reaction is first order with respect to peroxodisulphate and the rate is essentially independent of oxalate concentration (there is a slight increase in the first-order rate coefficient with increase of oxalate concentration). Allyl acetate inhibits the reaction and reduces the rate to that observed in the absence of oxalate. In the range pH 0.5-10.3 a rate maximum occurs at pH 4.5. The first-order rate coefficient for the reaction using 0.08 M disodium oxalate is expressed by... 

Early work showed that the rate of the silver ion-catalysed oxidation of oxalate is much faster than the oxidations of other substrates under similar conditions King ). Allen showed that with solutions of very low copper concentration, but not de-aerated, the rate is only slightly faster compared with other substrates. However, Kalb and Allen found that oxygen is a powerful inhibitor of the silver ion-catalysed oxidation, and that in the absence of oxygen low concentrations of copper have no effect on the rate. They studied the silver ion-catalysed reaction in the absence of oxygen. With peroxodisulphate concentrations greater than 0.004 M the rate equation is... 

Allyl acetate inhibits the reaction, but the maximally inhibited rate is 2.5 times faster than the silver ion-catalysed decomposition of peroxodisulphate in the absence of oxalate. With peroxodisulphate concentrations less than or equal to 0.004 M, the rate becomes proportional to the peroxodisulphate concentration squared and independent of the catalyst concentration, viz. 

Ben-Zvi and Allen studied the cupric ion-catalysed reaction. As in the uncatalysed reaction the kinetics are first order with respect to peroxodisulphate and zero order with respect to oxalate. The order with respect to total copper is half, the rate expression being... 

The reaction is approximately first-order with respect to each reactant (the second-order rate coefficient increases with increase of substrate concentration), and catalysis by hydroxide ions is observed. Henderson and Winkler studied the ferrous ion-catalysed oxidation of thioglycolicacidto dithioglycolic acid. The rate is sensitive to traces of metal ions, and reproducible results could not be obtained in the absence of the catalyst. The oxidation is first-order with respect to both peroxodisulphate and ferrous ions, and zero-order with respect to the substrate. The second-order rate coefficient is approximately equal to that determined in the absence of the substrate, so Henderson and Winkler suggested that the ratedetermining step is the oxidation of ferrous to ferric ions, as in reaction (96), and that this is followed by reaction (97) and then rapid oxidation of thioglycolic acid by ferric ions. 

The rate coefficient 2 is proportional to the concentration of phenolate ion thus for 2-hydroxy-pyridine (pA A = 11-6) oxidation below pH 8 is slower than the spontaneous decomposition of peroxodisulphate. For the reaction of 2-hydroxy-P5rridine in 2 M sodium hydroxide, the variation of 2 with temperature is expressed by... 

Metal ions are known to catalyse oxidations involving peroxodisulphate, especially silver(i), and the catalysis of the oxidation of [Ru(bipy)3] + has been described in the presence of this ion. The chelated complex is inert to hydrolysis and oxidised only slowly in the absence of metal ions. The mechanism in sulphuric acid (Scheme 11) involves the formation of a weak 1 1 complex between Ag+ and SaOg , with subsequent rapid reactions of... 

The oxidation kinetics of // fl i -aquabis(ethylenediamine)isothiocyanatocobalt(iii) ion with peroxodisulphate have been investigated in 0.01 M-HC104. The reaction products /rfl 5-[Co(OH2)(en)2CN]2+ and rran5-[Co(OH2)(en)2NH3]= + were identified by ion-exchange procedures. The rate law may be expressed as... 

Similar conclusions were reached in a second paper dealing with the pulse-radiolytic reduction of Cd + alone. Further reactions of these ions are the subject of another paper and, besides electron-transfer processes, additions of unsaturated compounds and oxygen atom transfers from NgO are described. The products of these latter reactions, MO+, behave like OH in their reactions with halide ions. In the electron-transfer processes, the reactions with peroxodisulphate proceed by a second-order pathway with rate constants of ca. 10 1 mol s ... 

The kinetics and mechanism of the reaction between peroxodisulphate and formate ions have been studied over the pH range 0.7—12.6. In three concentration ranges of [HC02"1, the rate law assumes three different forms. At [HCOa ]<0.01 mol 1 and [Sa08 "] 4x 10 molH, in the pH range 4.5—7 and in the absence of oxygen, the rate law is... 

Salt Effects.—Classical studies involving rate dependence on ionic strength as evidence of mechanism have included those of the reactions of c -[CoX2(en)2]+ with 1,10-phenanthroline, of c/5-[CoCl(NCS)(en)2] with periodate and with peroxodisulphate, and of formation of [AlF] +. Ion-pairing has been invoked to explain observed salt effects on the hydrolysis of 12-tungstosilicate (see Chapter 2 of this Part), on the base hydrolysis of (Co(NH3)5Cl] + in the presence of malonate, succinate, sulphate, or tripolyphosphate, and on the reaction between the [Fe(CN)6] and [W(CN)8] anions. However, in the reaction of [Fe(CN)6l with [8203] the marked effects of added alkali-metal cations were more easily explicable by their incorporation into the transition state ((NC)gFe]--M -[S2O3] than by initial-state ion-pairing. The observed acceleration of solvoly-... 

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