Kinetics of Acetate Decomposition

RATES This data block supplies a program written in basic which defines the rate reactions, using parameters specified in the KINETICS block. In addition to most of the usual basic commands, several commands unique to phreeqc are also available, which allow access to parameters calculated by phreeqc. Using a basic program means a bit more work for the user, but the results are worth it, in the sense that virtually any kinetic rate law may be introduced. 

3A recent release of The Geochemist s Workbench adds the ability to program any rate law, much as in this phreeqc example. 

Acetate Example Using phreeqc The acetate calculation can also be performed using phreeqc, using a simple basic program in the RATES data block as before. The script would be similar to that in Table 11.6. 

Part of almost all homogeneous kinetic calculations will be some method to decouple the reactive species, which are often redox species. In kinetic calculations, the species are obviously not at equilibrium with each other, at least at the start of the calculation they approach equilibrium during the calculation. But speciation programs such as phreeqc and react assume all species to be at equilibrium unless told otherwise. In this case we want CH4 to not react with other species, partly to see what it is doing during the reaction, and partly because in nature it is extremely unreactive. We also want acetate to be decoupled because it is metastable and will not even exist in the solution at equilibrium. At the time of writing, this is not necessary in phreeqc, because the 

In phreeqc, decoupling is achieved by defining new species and, if desired, their reactions with other species. In this case we simply define Methane to be a new species, and give it no reactions it is inert. In react (below), decoupling is achieved with a decouple statement. 

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Fig. 1. Kinetics of the decomposition of 0.2 mole of decanoyl peroxide in acetic acid and quinoline at 70°. ...

The kinetics of the decomposition of Tetralin and tert-butyl hydroperoxides by dilauryl thiodipropionate were determined in the temperature range 30°-90°C. The reaction had an induction period which was eliminated by acetic acid but not by dilauryl sulfinyldipropionate. The rate equation... 

Carothers WW, Kharaka YK (1978) Aliphatic acid anions in oil-field waters - implications for origin of natural gas. Am Assoc Pet Geol Bull 62 2441-2453 Child WC Jr, Hay AJ (1963) The thermodynamics of the thermal decomposition of acetic acid in the vapor phase. J Am Chem Soc 86 182-187 Clark LW (1957) The kinetics of the decomposition of oxalic acid in non-aqueous solutions. J Phys Chem 61 699-701... 

There have been numerous studies on the kinetics of decomposition of A IRK. AIBMe and other dialkyldiazenes.46 Solvent effects on are small by conventional standards but, nonetheless, significant. Data for AIBMe is presented in Table 3.3. The data come from a variety of sources and can be seen to increase in the series where the solvent is aliphatic < ester (including MMA) < aromatic (including styrene) < alcohol. There is a factor of two difference between kA in methanol and k< in ethyl acetate. The value of kA for AIBN is also reported to be higher in aromatic than in hydrocarbon solvents and to increase with the dielectric constant of the medium.31 79 80 Tlic kA of AIBMe and AIBN show no direct correlation with solvent viscosity (see also 3.3.1.1.3), which is consistent with the reaction being irreversible (Le. no cage return). 

Hisatsune and co-workers [290—299] have made extensive kinetic studies of the decomposition of various ions in alkali halide discs. Widths and frequencies of IR absorption bands are an indication of the extent to which a reactant ion forms a solid solution with the matrix halide. Sodium acetate was much less soluble in KBr than in KI but the activation energy for acetate breakdown in the latter matrix was the larger [297]. Shifts in frequency, indicating changes in symmetry, have been reported for oxalate [294] and formate [300] ions dispersed in KBr. 

There have been relatively few detailed kinetic studies of the decompositions of metal acetates, which usually react to yield [1046] either metal oxide and acetone or metal and acetic acid (+C02 + H2 + C). Copper(II) acetate resembles the formate in producing a volatile intermediate [copper(I) acetate] [152,1046,1047]. 

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]. 

When the decomposition of the zwitterionic intermediate is rate-determining, the effect of the solvent is crucial since it may produce changes in the mechanisms and in the rate-determining step. A recent study of the kinetics of the reactions of 1-chloro-, 1-fluoro- and l-phenoxy-2,4-dinitrobenzene with piperidine, n-butylamine and benzylamine in ethyl acetate and THF indicated that these reactions resemble those in dipolar aprotic solvents when primary amines are the nucleophiles (i.e. that shown in equation 1, with... 

The kinetic and activation parameters for the decomposition of dimethylphenylsilyl hydrotrioxide involve large negative activation entropies, a significant substituent effect on the decomposition in ethyl acetate, dependence of the decomposition rate on the solvent polarity (acetone-rfe > methyl acetate > dimethyl ether) and no measurable effect of the radical inhibitor on the rate of decomposition. These features indicate the importance of polar decomposition pathways. Some of the mechanistic possibilities involving solvated dimeric 71 and/or polymeric hydrogen-bonded forms of the hydrotrioxide are shown in Scheme 18. 

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. 

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. 

The kinetics of the noncatalytic reaction were studied by the method of Bawn and Williamson (4). They found 3.3-3.7 moles/liter for the equilibrium constant for the formation of the intermediate acetaldehyde monoperacetate (AMP) and a first-order rate constant for the decomposition of this intermediate to acetic acid of 0.015 min."1 at 25°C. We found difficulty in reproducing our results probably caused mainly by the high values of the blanks in the iodometric methods used. However, as an average of four determinations we obtained 0.03 min."1 at 30°C. 

The kinetics of formation of ketene and acetic acid on thermal unimolecular decomposition of acetic anhydride at 750-980 K have been reported and used to reevaluate the Arrhenius equation as k = 10122exp(—145 kJ mol l/RT) s-1 for the temperature range 470-980 K.54 Results of ah initio MO calculations suggest that the reaction proceeds by concerted elimination through a six-centre transition state, with potential barrier height 156 kJ mol-1. 

Substrate-catalyst interaction is also essential for micellar catalysis, the principles of which have long been established and consistently described in detail [63-66]. The main feature of micellar catalysis is the ability of reacting species to concentrate inside micelles, which leads to a considerable acceleration of the reaction. The same principle may apply for polymer systems. An interesting way to concentrate the substrate inside polymer catalysts is the use of cross-linked amphiphilic polymer latexes [67-69]. Liu et al. [67] synthesized a histidine-containing resin which was active in hydrolysis of p-nitrophenyl acetate (NPA). The kinetics curve of NPA decomposition in the presence of the resin was of Michaelis-Menten type, indicating that the catalytic act was accompanied by sorption of the substrate. However, no discussion of the possible sorption mechanisms (i.e., sorption by the interfaces or by the core of the resin beads) was presented. 

Kinetic Order of the Reaction. The decomposition of the hydroperoxide had an initial slow reaction or induction period followed by a faster main reaction. The induction period was unaffected by the addition of dilauryl sulfinyl dipropionate or by carrying out the reaction in an atmosphere of nitrogen but was eliminated by the addition of acetic acid. The length of the induction period decreased as the initial concentration of both hydroperoxide or sulfur compound increased. 

The thermal decomposition of metal sulphates, acetates, oxalates, nitrates, and so forth can also be considered vdth similar thermodynamic considerations. Because the partial pressure of these gaseous decomposition products are minuscule in air, these salts are unstable at all temperatures where AG is negative. The kinetics of decomposition of these salts, however, is slow at low temperatures. These reaction thermodynamic fundamentals are applicable to all other reactions discussed in this chapter. 

The kinetics of reactions on metal surfaces is strongly affected by the structures of reaction intermediates, especially by the incorporated metal adatoms in their structures (Sect 11.5). hi the mid 1970s, Falconer and Madix observed a surface- kinetic explosion for the decomposition of formate and acetate adsorbed on the Ni (110) surface [23, 24], Recently, with the help of STM, TPRS, and XPS, we were able to determine that Ni atoms are incorporated into the structures of the carboxylate intermediates. Remarkably, the incorporation of metal atoms into the carboxylate structure is an important aspect of the origin of the kinetic explosion. 

In the mid 1970s, Falconer and Madix observed a surface- kinetic explosion for the decomposition of formic acid (HCOOH) [23] and acetic acid (CH COOH) [24] on the Ni(llO) surface, characterized by very narrow product desorption peaks in TPRS. Such autocatalytic reactions have also been observed in the decomposition of acetic acid on Pd(llO), Rh(llO), Rh(lll), and even supported Rh catalyst by Bowker et al. [70-75]. In general, these reactions exhibit accelerations in rate as the reaction proceeds to completion. Earlier work hypothesized that decomposition of the carboxylate species formed following adsorption of the acids on the surface was initiated at vacancies (i.e. bare metal sites) and propagated by the further creation of vacancies as the products desorbed from the surface [23, 24]. The rate of decomposition was well described by the rate equation r = -k(C / Cj )(Cj - c+/Cj), in which C is the instantaneous surface concentration of carboxylate, C, is the initial surface concentration, and/is the density of initiation sites. Since the decomposition produced an ever-increasing concentration of vacant sites, a kinetic explosion occurred. 

Annealing the formate-covered Ni(llO) surface to 340 K, the leading edge of formate decomposition produces scattered Ni islands of monatomic step height amidst the c(2 x 2)-formate domains and the small number of pits of monatomic step depth (Fig. 11.15) [21]. The Ni islands are apparently formed by Ni adatoms released during the decomposition of a small fraction of formate. The Ni islands were also observed when annealing the acetate-covered Ni(llO) surface to 360 K [21]. These observations provide a more detailed explanation for the kinetic explosion in the decomposition of formate and acetate. It appears that the released Ni atoms catalyzed the further decomposition of the carboxylates. The released Ni atoms can either catalyze the reaction before they nucleate into islands or provide unoccupied metal sites, which are produced exponentially due to desorption of the carboxylates with the exponential rate, for decomposition. This ultimately leads to a kinetic explosion in the decomposition of the carboxylates on Ni(llO). 

Mohamed et al. [42] have reported non-isothermal kinetic studies of the decompositions of nickel and lead acetate hydrates. They review previous studies and report analyses of reaction products. 

Comparisons of observations for the decompositions of five metal malonates (Cu(II) [22], Ag [62], Ni [65], Co [66] and Ca [67]) illustrate the behavioural idiosyncrasies and individualities of these closely related compounds. Anion breakdown was accompanied by hydrogen transfer so that four of the reactants gave appreciable yields of acetate or acetic acid. (No analytical measurement was made for the cobalt salt.) The principal kinetic results (Table 18.3.) show the remarkably contrasting behaviours of copper(II) malonate, where decomposition... 

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