Oxalic acid activation energy

The kinetics of the Ce(IV) sulphate oxidation of oxalic acid are simple second order although the rate coefficient is inversely proportional to both hydrogen and bisulphate-ion concentrations, and it is also reduced at very high oxalic acid concentrations. Values of the activation energy from 13.4+1.5 (ref. 411) to 16.5+0.4 (ref. 409) kcal.mole have been reported. An intermediate has been detected spectroscopically " this decays in first-order fashion with E = 10.5+0.5 kcal.mole" and with a rate independent of acidity. However, the extent of formation of this complex is reduced as the acidity is increased ", and it appears that a less reactive dioxalato complex is formed at higher substrate concentrations ". 

Boldyrev et al. [46], from quantum mechanical calculations of bond strengths in the oxalate anion, and from observations [38] of the decomposition of this species in potassium bromide matrices, concluded that the most probable controlling step in the breakdown of the oxalate ion is rupture of the C-C bond. This model is (again) based on the observation that the magnitudes of the activation energies for decompositions of many metal salts of oxalic acid are comparable. This model was successfiilly applied [46,68] to the decompositions of many oxalates, with the possible exception of silver oxalate where the strengths of the C-C and Ag-0 bonds are similar. 

The decomposition of oxalic acid photosensitized by uranyl ion is a common actino-metric reaction. The light may have any wavelength between 250 and 450 nm. The absorption of the light quantum activates the uranyl ion, which transfers its energy to a molecule of oxalic acid, which then decomposes. The reaction may be written... 

In the decomposition of oxalic acid (39) to yield CO2 and formic acid, the energy of activation was constant in dioxane-water mixtures, but decreased with increasing temperature in aqueous solution. This was ascribed to the activated complex having fewer degrees of freedom than the reactant. 

Clearly, regime 1 is the simplest to analyze because it is kinetically controlled and therefore for all practical purposes can be treated as a homogeneous system. Even here, however, the solubility of A in phase 2 (the liquid), its partial pressure, and solvent vapor pressure can exert a significant influence. Thus, although the rate constant increases with increase in temperature, the solubility generally decreases, but the overall effect would still be positive because the activation energy is almost always higher than the heat of solution. On the other hand, the solubility can also increase with temperature (particularly for liquids), for example, chlorobenzene in aqueous sulfuric acid and esters such as ethyl p-nitrobenzoate and dichloroethyl oxalate in water. This serves only to supplement the effect of temperature on the rate constant. 

With a high acid concentration in the electrolyte, thermal precipitation of VOj can be prevented because VOj is stable due to the dimerization of VOj irais to V204 and 203" " species. However, with an increased acid concentration, the solubility of V(II)W(III) species decreases, but the viscosity increases obviously. This means that more energy is consumed to pump the electrolyte into the battery. The concentration of acid should be in the range of 2-4 M. Recently, to increase electrochemical activity and the concentration of vanadium ions in the electrolyte, various organic additives, such as sodium hexametaphosphate, alkali metal sulfate, and alkali metal oxalate, have been investigated and added into the RFBs electrolytes [21, 22]. 

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