Acetic acid reaction enthalpy

One of the issues of the industrial process design is related to the heat released by this reaction. A temperature rise will decrease the acetic acid yield, not only because the equilibrium constant becomes lower (the reaction is exothermic see section 2.9) but also because it will reduce the enzyme activity. It is therefore important to keep the reaction temperature within a certain range, for instance, by using a heat exchanger. However, to design this device we need to know the reaction enthalpy under the experimental conditions, and this quantity cannot be easily found in the chemical literature. 

Let us suppose that the acetic acid content of the final aqueous solution is 5%, corresponding to a ratio of approximately 1 mol of CH3COOH to 60 mol of H2O. As the yield of reaction 2.1 will be near 100% (recall that reaction 2.2 is rather exothermic, implying a very high equilibrium constant see section 2.9), the same value will be used for the molar ratio (H2 O) / n (C 2115OII), despite the increased total amount of substance of water in the reaction products. In the present case, the difference of 1 mol of water between the product and the reactant mixtures has a negligible enthalpic effect. The enthalpies associated with the solution of ethanol and acetic acid in 60 mol of water are derived from literature data [17] as Asin//(1) = -10.0 0.1 kJ mol-1 and Asin//(3) = —1.0 0.1 kJ mol-1. This calculation will be detailed in section 2.5. 

In contrast to the acetic acid case study, there are many important reactions wherein the temperature has a significant effect on the enthalpies. Consider, for example, reaction 2.13, which represents one of the industrial processes for manufacturing acetylene using methane (from natural gas) as the starting material. 

Why did we prefer to use the cycle in figure 2.1 instead of the easier method after equation 2.21 Simply because we had considered that the best data for the standard enthalpies of formation of pure ethanol and acetic acid are those recommended in Pedley s tables [15]. The values (-277.6 kJ mol-1 and -484.3 kJ mol-1, respectively) are both about 1 kJ mol-1 less negative than those in the NBS Tables, and their difference nearly cancels when the reaction enthalpy is calculated. But of course, we are seldom so lucky. Using data from different databases may lead to much larger discrepancies. 

As appears from the examination of the equations (giving the best fit to the rate data) in Table 21, no relation between the form of the kinetic equation and the type of catalyst can be found. It seems likely that the equations are really semi-empirical expressions and it is risky to draw any conclusion about the actual reaction mechanism from the kinetic model. In spite of the formalism of the reported studies, two observations should be mentioned. Maatman et al. [410] calculated from the rate coefficients for the esterification of acetic acid with 1-propanol on silica gel, the site density of the catalyst using a method reported previously [418]. They found a relatively high site density, which justifies the identification of active sites of silica gel with the surface silanol groups made by Fricke and Alpeter [411]. The same authors [411] also estimated the values of the standard enthalpy and entropy changes on adsorption of propanol from kinetic data from the relatively low values they presume that propanol is weakly adsorbed on the surface, retaining much of the character of the liquid alcohol. 

Two excellent reviews <71AHC(13)235, 72IJS(C)(7)6l) have dealt with quantitative aspects of electrophilic substitution on thiophenes. Electrophilic substitution in the thiophene ring appears to proceed in most cases by a mechanism similar to that for the homocyclic benzene substrates. The first step involves the formation of a cr-complex, which is rate determining in most reactions in a few cases the decomposition of this intermediate may be rate determining. Evidence for the similarity of mechanism in the thiophene and benzene series stems from detailed kinetic studies. Thus in protodetritiation of thiophene derivatives in aqueous sulfuric and perchloric acids, a linear correlation between log k and —Ho has been established the slopes are very close to those reported for hydrogen exchanges in benzene derivatives. Likewise, the kinetic profile of the reaction of thiophene derivatives with bromine in acetic acid in the dark is the same as for bromination of benzene derivatives. The activation enthalpies and entropies for bromination of thiophene and mesitylene are very similar. 

Consider the following example. AG° for dissociation of acetic acid in water is +27.2 kj/mol (Table 6-5). The enthalpy change AH0 for this process is almost zero (-0.1 kj/mol) and AS0 is consequently -91.6 J I< This large entropy decrease reflects the increased amount of water that is immobilized in the hydration spheres of the H+ and acetate- ions formed in the dissociation reaction. In contrast, dissociation of NH4+ to NH3 and H+ converts one positive ion to another. AH0 is large (+52.5 kj/mol) but the entropy change AS0 is small (-2.0 J K-1, Table 6-5). 

Use the information in Appendix 2A to determine whether this reaction is exothermic or endothermic by calculating the standard reaction enthalpy, (b) Acetic acid can also be formed by the oxidation of ethanol ... 

In perchloric acid, hexoses and pentoses are oxidized by Ce(IV) via formation of two complex intermediates. The first is partly oxidized following Michaelis-Menten kinetics and partly dissociated to the second, which is oxidized more slowly than the former.180 The first step in the oxidation of aldoses by Tl(III) in the same medium involves the C-l-C-2 cleavage of the aldehydo form of the sugar. Thus, D-glucose gives D-arabinose and formic acid. With an excess of oxidant the final product is carbon dioxide.181 In the presence of a catalytic amount of sulfuric acid in acetic acid, Tl(III) oxidizes maltose and lactose to the corresponding disaccharide aldonic acids. The reaction showed activation enthalpies and enthropies characteristic of second-order reactions.182... 

The enthalpy changes for adsorption of acetaldehyde (step 3), ethanol (step 5), hydrogen (step 6), water (step 8), and acetic acid to form adsorbed acetate (step 9) were adjusted in the reaction kinetics analysis. The initial estimates of the heats of adsorption of acetaldehyde, ethanol, and hydrogen were obtained from the DFT predictions for these species on Cu(211) (Table VIII). The heat of adsorption of water was constrained to be equal to the heat of adsorption of ethanol in these analyses. The steps involving adsorption of ethanol, acetaldehyde, water, and the step in which acetic acid forms the surface acetate species were all assumed to be nonactivated. 

The standard entropy and enthalpy changes for the dissociative adsorption of ethyl acetate (step 1), acetic acid (step 7), and the hydrogenation of adsorbed acetaldehyde to form ethoxy species (step 4) were constrained to maintain thermodynamic consistency for the appropriate overall gas-phase reactions. 

Most of the uncertainty in these enthalpies of formation for the chlorinated acetyl chlorides, RCOC1, arise from uncertainties in the corresponding chlorinated acetic acids, RCOOH after all, it is the high-accuracy (basic, aqueous) hydrolysis reaction that interconnects these species that gives us the acyl chloride enthalpies of formation we use. [See G. M. Moselhy and H. O. Pritchard, J. Chem. Thermodyn., 7,977 (1975).] We note that sufficiently few chlorinated acetate esters enjoy sufficiently accurate enthalpies of formation (as determined by combustion measurements) to allow for comparison of RCOC1 and RCOOR for any R. ... 

The enthalpy of reaction, AH, is the other important thermodynamic parameter to consider. On its own, whether a reaction is exothermic or endothermic will not determine if a reaction is industrially feasible or not. Both exothermic and endothermic processes are known in industry, methanol carbonylation to acetic acid (Equation 3 AH —123 kJ/mol at 200°C), being an example of the former and the steam reforming of methane to synthesis gas, (Equation 4 AH + 227 kJ/mol at 800°C), being an example of the latter. 

The observed rate coefficient, k, for the acid-catalysed proton exchange has a value of 1.46 x 107 1 mole 1 sec-1, the reaction having an enthalpy of activation of 5.6 kcal mole"1. As found for the analogous reaction of acetic acid previously described, the observed rate is much faster than the predicted rate of proton loss from the protonated phenol, so that proto-nated phenol is not formed as an intermediate. The reaction proceeds by the concerted mechanism... 

The kinetics of the optical rotatory changes of poly-L-proline in various solvents and at various temperatures have also been studied by Steinberg et al. (1960a). In acetic acid the course of the forward mutarotation reaction was found to be independent of concentration (over the range 0.25 to 2.0 gm/KK) ml) but, as observed by Downie and Randall, the rate constant depends on the degree of mutarotation. An activation enthalpy, AH = 21 kcal/mole, was determined for both the forward mutarotation of form I in acetic acid and the reverse mutarotation of form II in acetic acid-w-pro-panol. 

The type of mutarotation kinetics found experimentally is strongly dependent on the solvent. Forward mutarotation in acetic acid yields a plot of log d[a /dt versus log ([a]( — [ ] ) which is linear over 97% of the reaction with a slope of 1.33. On the other hand the rate is virtually invariant over two-thirds of the reaction in a solvent of 30% water-70 % acetic acid. Reverse mutarotation in acetic acid-n-propanol appears to follow first-order kinetics over 90% of the transformation. From these experiments and the fact that the enthalpies of form I and form II are essentially identical (Steinberg et al., 1960a) it seems likely that solvation is decisive not only in determining the kinetic pathway of mutarotation but also the structural pattern which is stabilized in a given solution. 

Use the data in Table 6-9 to calculate the enthalpy and entropy of reaction for dissociation of acetic acid using (a) Equations (4) and (5) and (b) the temperature dependence of of Equation (7) by graphing In Kg versus 1/T. 

Equation 21 represents an acetylation reaction of an aniline, where both the acetic acid and water are liquids and the nitrogen-containing species are solids. For Ar = Ph, the enthalpy of reaction is —32.7 kJmol-1. 

To understand of specific phenomena on the basis of a few simple principles. Examples (i) the concept of the equilibrium constant for dissociation of acetic acid, (ii) the effect of surface potential on the surface reaction enthalpy. 

To provide the relationships which enable calculation of a certain property (physical quantity) from other measured data. Examples (i) evaluation of the equilibrium constant of the dissociation of acetic acid from pH measurements, (ii) evaluation of surface reaction enthalpy from the temperature jump in a calorimeter, or from the temperature dependency of the point of zero charge. 

Enthalpies of vaporization are 35.23 (methanol) and 24.31 (acetic acid) kJ/mol. Using data from Appendix III estimate AH for this reaction. 

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