Reactions of peroxodisulphate

Reviews of the mechanisms of many reactions of peroxodisulphate have been published They provide fuller discussion of many points raised here. 

Extensive publication on the thermal decomposition reaction makes this one of the most fully studied systems in the whole of chemistry, even after 

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 

The peroxomonosulphuric acid so formed is supposed to decompose rapidly to form oxygen at acid concentrations less than or equal to 0.5 M under the experimental conditions used. The result of the oxygen-tracer experiment (O2 from S20 ) resembles observations of hydrogen peroxide decompositions. The reaction does not show inhibition related to the concentration product [H ] 

The mechanism of reaction is thought to involve protonation of peroxodisulphate (4). An early mechanism, written in terms of sulphur tetroxide (5), is almost certainly incorrect. The sequence (6)-(8), avoids 

Oxidation rates involving this anion are usually slow and in many instances the rate-determining step is the homolytic cleavage of the 0—0 bond. The presence of a metal cation may increase the rate although the mode of reaction and the rate law vary with the nature of M +. The copper(n)-catalysed oxidation of succinic acid proceeds according to the rate law  

A Cu -succinate (CuLa) complex is postulated which may undergo oxidation to a transient Cu species, e.g, 

The hydrogen ion dependence is attributed to the inhibition of formation of the copper complex but the mechanism will hold only if [L] [Cu ]. A similar halforder dependence on Cu is exhibited in the corresponding oxidation of glycollic acid. In this instance, however, the copper complex [Cu(gly)al is involved directly in a redox reaction with the oxidant  

A feature of both systems is the autocatalysis by SO4 and COa (derived from the reductant) radicals. A simple rate law is observed in the Ag -catalysed oxidation of aspartic acid (Rate = A a[S208 ][Ag ][Asp] ). Medium effects are important with inhibition by ions in the order Mg + K+ Na+ H+ and NOs S04 . The uncatalysed oxidations of glyoxal and glyoxylic acid have also been investigated. Potassium ruthenate (K2RUO4), which can be readily prepared from reaction of ruthenium trichloride with aqueous persulphate, can be used cata-lytically in the presence of 8208 for the oxidation of organic substrates under mild conditions. RuO is considered to act as a two-electron oxidant. 

Catalysis by cobalt(ii) has been observed in the oxidation of iodide by per-oxodisulphate, the mechanism involving formation of the reactant complex C0I+. It is perhaps of interest that this is the first paper concerned with a redox mechanism that the Reporter has seen from this source. 

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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. 

Detonation occurs during the reaction of molten aluminium with ammonium peroxodisulphate in the presence of water. However, since the temperature is above 75°C, the presence of water is sufficient to decompose it and the water/molten aluminium interaction has aiready been mentioned as being explosive. 

Schuchmann H-P, von Sonntag C (1988) The oxidation of methanol and 2-propanol by potassium peroxodisulphate in aqueous solution free-radical chain mechanisms elucidated by radiation-chemical techniques. Radiat Phys Chem 32 149-156 Schwarz HA, Bielski BHJ (1986) Reactions of H02 and 02 with iodine and bromine and l2 and I atom reduction potentials. J Phys Chem 90 1445-1448... 

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 ... 

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)... 

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... 

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. 

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. 

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... 

Figure 4 Correlation between free-energy of activation and the standard free-energy of the redox step for reactions of (1) [Fe(4,4 -dimethyl-bipy)3] + (2) [Fe(bipy)3] + (3) [Fe(5-Me-phen)g]2+ (4) [Fe(phen)3] + with peroxodisulphate...

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 reaction of azide with peroxodisulphate in neutral or acidic solutions proceeds with the stoicheiometric amounts of dinitrogen being formed ... 

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... 

The rate law in the reduction of peroxodisulphate by [Fe(CN)6S03] - indicates two reaction pathways ... 

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-... 

The rate-determining step in the reaction of [Fe(bipy)3] +, [Fe(phen)3] +, or [Fe(terpy)2] + with peroxodiphosphate is dissociative - the breaking of an iron-nitrogen bond. It has been recognized for some time that these and related iron(n) complexes react with peroxodisulphate by parallel oxidative and dissociative rate-determining steps the reaction of the [Fe(ppsa)3] + cation mentioned above with peroxodisulphate provides a recent example of this behaviour. It is now apparent that any direct oxidation of this type of iron(n) complex with peroxodiphosphate is so slow that reaction proceeds entirely via rate-determining dissociation, presumably with subsequent rapid oxidation of the products of dissociation. 

Potassium peroxodisulphate (K2S2Og) also oxidizes sulphoxides to sulphones in high yield, either by catalysis with silver(I) or copper(II) salts at room temperature85 or in pH 8 buffer at 60-80 °c86-88. The latter conditions have been the subject of a kinetic study, and of the five mechanisms suggested, one has been shown to fit the experimental data best. Thus, the reaction involves the heterolytic cleavage of the peroxodisulphate to sulphur... 

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

However, although the oxidation process was ascribed to either direct reaction on the electrode surface, or mediated by peroxodisulphate and other inorganic reagents electrogenerated at the anode surface, the linear decrease of the Faradic yield down to zero in the investigated range of concentration was interpreted as an indication of a process under diffusive control. This leads to the conclusion that the oxidative degradation of the compounds essentially occurred at the electrode interface. 

To study these reactions use a freshly prepared 0 1m solution of ammonium peroxodisulphate (NH4)2S208. 

It must be emphasized that these reactions take place only in acid media. Add solid potassium peroxodisulphate and dilute sulphuric acid to the neutralized soda extract of the mixed halides contained in a conical flask heat the flask to about 80°, and aspirate a current of air through the solution with the aid of a filter pump until the solution is colourless (Fig. IV.2. T is a drawn out capillary... 

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