Peracetic decomposition

EFFECT OF THE ALKYL GROUP (R) ON THE RATE OF PERACETATE DECOMPOSITION ... 

Little work has been carried out on thiazole N-oxides. These products are unstable and breakdown by autoxidation to give thiazolium-A -oxide sulfates and other decomposition products (264). They are prepared by direct oxidation with hydrogen peroxide, or by tungstic acid (264, 265) or peracetic acid (265-267). 

Dibromoacetic acid [631-64-1] (Br2CHCOOH), mol wt 217.8, C2H2Br202, mp 48°C, bp 232—234°C (decomposition), is soluble in water and ethyl alcohol. It is prepared by adding bromine to boiling acetic acid, or by oxidi2ing tribromoethene [598-16-3] with peracetic acid. 

In the case of reaction 3, entries 1 and 2, that is, iert-butyl peracetate and (ert-butyl perpropionate, almost certainly decompose by a stepwise mechanism, rather than the concerted mechanism assumed for reaction 3. Entry 3, tert-butyl perisobutyrate, probably forms the least stable R radical by the perester decomposition mechanism which is still mostly concerted in nature (36). 

The only accident that involves a saturated ester is the result of an attempt to extract an organic residue containing hydrogen peroxide with ethyl acetate. The latter was mixed with methanol and refluxed with the residue and hydrogen peroxide in an aqueous solution. A second extraction was carried out with acetate and the liquid was then evaporated. The small quantity of the compound that remained after the evaporation detonated violently. It was thought that this detonation was the result of the violent decomposition of methyl hydroperoxide, peracetic acid and/or ethyl peracetate. 

The thermal decomposition of peracetic acid in an aqueous solution produces acetic acid, C02, and dioxygen [4]. The detailed data on the chemistry and the kinetics of peracids decay are presented elsewhere [4—7]. 

A 500 ml bottle of peracetic acid (ethaneperoxoic acid), provided with a vented cap, was received packed within a tin padded with what appeared to be paper. The metal was distinctly hot to the touch on opening. Peroxyacetic acid is capable of runaway decomposition from 65°C. The heat is thought to have resulted from... 

Aliphatic amines have much less effect on the later reactions of the gas-phase oxidation of acetaldehyde and ethyl ether than if added at the start of reaction. There is no evidence that they catalyze decomposition of peroxides, but they appear to retard decomposition of peracetic acid. Amines have no marked effect on the rate of decomposition of tert-butyl peroxide and ethyl tert-butyl peroxide. The nature of products formed from the peroxides is not altered by the amine, but product distribution is changed. Rate constants at 153°C. for the reaction between methyl radicals and amines are calculated for a number of primary, secondary, and tertiary amines and are compared with the effectiveness of the amine as an inhibitor of gas-phase oxidation reactions. 

Similar results are obtained if the inhibitor is added at a later stage (when more peracetic acid is present), but the induction periods are even shorter (Figure 3). Indeed, at the minimum pressure change, when the maximum concentration of peracetic acid is present, it is only when a large amount of diethylamine has been added (over 10% of the total mixture) that decomposition appears to stop (Figure 4). 

The results given in this paper show that aliphatic amines do not catalyze the decomposition of peroxides, and compared with their effect at the start of reaction, they have much less effect on the later stages of oxidation, although they appear to retard the decomposition of peracetic acid. The reactions of radicals with aliphatic amines indicate that an important mode of inhibition is most probably by stabilization of free radicals by amine molecules early in the chain mechanism, possibly radicals formed from the initiation reaction between the fuel and oxygen. For inhibition to be effective, the amine radical must not take any further part in the chain reaction set up in the fuel-oxygen system. The fate of the inhibitor molecules is being elucidated at present. 

Ethyl peracetate was the first ester of a peroxy acid, and was characterized by Baeyer and Villiger in 1901. Kinetic studies of perester decomposition were reported by Blomquist and Ferris in 1951, and in 1958 Bartlett and Hiatt proposed that concerted multiple bond scission of peresters could occur when stabilized radicals were formed (equation 46). As noted below (equation 57), polar effects in perester decomposition are also significant. 

The oxidation of azo compounds with hydrogen peroxide in an acetic acid medium or with peracetic acid has been carried out by many investigators. For example, in a study of the 4,4 -dialkoxyazoxybenzenes, which are of interest in the field of liquid crystals, the corresponding alkoxynitrobenzenes were reduced with lithium aluminum hydride to the corresponding azo stage and, after decomposition of the reducing agent and removal of the solvent, the product residue was taken up in acetic acid and oxidized at 65°C for 36 hr with... 

Laundry powders in Europe use sodium percarbonate and a bleach activator, tetraacetylethylenediamine ( ED), to form peracetic acid in the washing machine that then reacts with bleachable stains such as tea, red wine, coffee, and curry to remove /bleach them from cloth. Transition metal ions, in particular copper, cause first decomposition of peracetic add before it can bleach stains and second react with peroxide to form highly readive hydroxyl radicals that can cause dye and fabric damage [32]. Chelants bind with copper and other metals to prevent these unwanted side readions. 

Mn 0.11, Cu 1.1, Zn 2.3 ppm. The residual peracetic acid after 30 min [37] is shown in Table 10.6. EDDS gave the most stable peracid release, followed by DTPMP and EDTMP. HEDP and other biodegradable chelants did not prevent peracid decomposition and there was no peracid remaining after 10 min under these conditions. 

To understand the significant effect of catalyst nature, a better understanding of the main reactions, peracetic acid decomposition, and its reaction with acetaldehyde was needed. A literature -survey showed that the kinetics were not well studied, most of the work being done at very low catalyst concentration 1 p.p.m.), and there is disagreement with respect to the kinetic expressions reported by different authors. The emphasis has always been on the kinetics but not on the products obtained, which are frequently assumed to be only acetic acid and oxygen. Consequently, the effectiveness of a catalyst was measured only by the rates and not by the significant amount of by-products that can be produced. We have studied the kinetics of these reactions, supplemented by by-product studies and experiments with 14C-tagged acetaldehyde and acetic acid to arrive at a reaction scheme which allows us to explain the difference in behavior of the different metal ions. 

Kinetic Studies. Peracetic Ac id Decomposition. Studies with manganese catalyst were conducted by the capacity-flow method described by Caldin (9). The reactor consisted of a glass tube (5 inches long X 2 inches o.d.), a small centrifugal pump (for stirring by circulation), and a coil for temperature control (usually 1°C.) total liquid volume was 550 ml. Standardized peracetic acid solutions in acetic acid (0.1-0.4M) and catalyst solutions also in acetic acid were metered into the reactor with separate positive displacement pumps. Samples were quenched with aqueous potassium iodide. The liberated iodine was titrated with thiosulfate. Peracetic acid decomposition rates were calculated from the feed rate and the difference between peracetic acid concentration in the feed and exit streams. 

Peracetic acid decomposition kinetics in the presence of cobalt or copper acetates were studied in the same apparatus used for the manganese-catalyzed reaction. However, in these studies it was used as a batch reaction system. The reactor was charged with peracetic acid (ca. 0.5M in acetic acid) and allowed to reach the desired temperature. At this time the catalyst (in acetic acid) was added. Samples were withdrawn and quenched with potassium iodide at measured time intervals. 

Peracetic Acid Decomposition. Peracetic acid solutions in acetic acid are stable at 30°C. in the absence of catalyst. Addition of copper (II)... 

Since copper (II) does not significantly catalyze the peracetic acid decomposition, we have studied the kinetics of this reaction only in the presence of manganese and cobalt acetates. 

With cobalt as catalyst the plot of log [peracetic acid] vs. time was linear for each cobalt acetate concentration. The first-order rate constants obtained at different cobalt concentrations (k2 ) were plotted as a function of total cobalt (Cot) concentration, and the plot indicates a first-order dependence on total cobalt as shown in Figure 3. The experimental rate law for the cobalt-catalyzed decomposition is thus ... 

Figure 1. Manganese-catalyzed decomposition of peracetic acid at 30°C.

Figure 2. Effect of manganese acetate on peracetic acid decomposition at...

Table IV. Effect of Temperature on the Metal-Catalyzed Peracetic Acid Decomposition...

Peracetic Acid Decomposition. Although, by comparing the rate of peracetic acid decomposition with the rate of its reaction with acetaldehyde, we can rule out the decomposition as a major path in acetaldehyde oxidation (see below), we will discuss the possible mechanisms for the catalytic decomposition of peracetic acid. 

Intermediate II has been postulated by Koubek et ah (18) for the bi-molecular decomposition of peracetic acid in aqueous solutions. 

An alternate mechanism for the cobalt (III)-catalyzed peracetic acid decomposition with a sequence of oxidation-reduction reactions, could be as follows ... 

With manganese and cobalt acetate the reaction of peracetic acid with acetaldehyde is very fast, and AMP is not detected. By comparing our rates with literature values of k17, k.17, and k18 we cannot propose a mechanism in which the only role of the metal ion is to catalyze the decomposition of AMP. The experimental rates in the presence of either manganese or cobalt acetates are much faster than the noncatalytic rate of formation of AMP. Thus, AMP per se is probably not an intermediate in the presence of these catalysts. 

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