Water-methanol equilibrium constant

Figure 29. Graph of the Logarithm of the Water/Methanol Association Equilibrium Constant against the Reciprocal of the Absolute Temperature...

Taking known values for the molar refractivities of water and methanol, and again assuming a range of values for the equilibrium constant (k) and the refractive index (ni) of the methanol/water associate, the actual values that fit the equation for these... 

It is seen that the three values for the equilibrium constant (k) range from 0.00443 to 0.00565 with an average value of 0.00504. The two values for the densities of the methanol/water associate are in reasonable agreement and have a magnitude that would be expected for the hydrogen bonded associate. 

Using the average value for the equilibrium constant, the distribution concentration of the different components of a methanol water mixture were calculated for initial methanol concentrations ranging from zero to 100%v/v. The curves they obtained are shown in Figure 28. The molar refractivities of 11.88 is also in accordance with that expected since the molar refractivity s of water and methanol are 3.72 and 8.28 respectively. The refractive index of the associate of 1.3502 is, as would be expected, higher than that of either water or methanol. 

Table 1 Equilibrium constants log K for the protonation of macrocyclic diamines [4] in water and methanol at 25°C.

A value of kjkp = 17 000 has been determined for partitioning of the acetophenone oxocarbenium ion [12+] in water.15,16 It is not possible to estimate an equilibrium constant for the addition of water to [12+], because of the instability of the hemiketal product of this reaction. However, kinetic and thermodynamic parameters have been determined for the reaction of [12+] with methanol to form protonated acetophenone dimethyl ketal [12]-OMeH+ and for loss of a proton to form a-methoxystyrene [13] in water (Scheme 10).15,16 Substitution of these rate and equilibrium constants into equation (3) gives values of AMeoH = 6.5 kcal mol-1 and Ap = 13.8 kcal mol-1 for the intrinsic... 

Another type of ternary electrolyte system consists of two solvents and one salt, such as methanol-water-NaBr. Vapor-liquid equilibrium of such mixed solvent electrolyte systems has never been studied with a thermodynamic model that takes into account the presence of salts explicitly. However, it should be recognized that the interaction parameters of solvent-salt binary systems are functions of the mixed solvent dielectric constant since the ion-molecular electrostatic interaction energies, gma and gmc, depend on the reciprocal of the dielectric constant of the solvent (Robinson and Stokes, (13)). Pure component parameters, such as gmm and gca, are not functions of dielectric constant. Results of data correlation on vapor-liquid equilibrium of methanol-water-NaBr and methanol-water-LiCl at 298.15°K are shown in Tables 9 and 10. 

This crude analysis is based on the behavior postulated by the Born equation. However, ion-pair formation equilibrium constants have been observed to deviate ma edly from that behavior (22/ -222)1 Oakenful, and Fenwick (222) found a maximum in the ion-pair formation constants of several alkylamines with carboxylic acids when determined at various methanol-water solvent compositions as shown by their data in Fig. 54. The results demonstrate that in this system the stability constant decreases with increasing organic solvent concentration above a.critical value which yields maximum stability. The authors suggested that this was due to a weakening of hydrophobic interactions between the ion-pair forming species by increased alcohol concentrations. In practice the effect of added organic solvent has been either to decrease the retention factor or to have virtually no effect. 

A notable difference between reactions of carbocations with water as a nucleophile and a base is the significantly higher intrinsic barrier for the latter. This difference has been demonstrated most explicitly by Richard, Williams, and Amyes22,246 for reaction of the a-methoxyphenethyl cation 68 with methanol (rather than water) acting as the base and nucleophile. The two reactions and the intrinsic barriers calculated from their rate and equilibrium constants are shown in Scheme 32. Values of A = 6.8 and 13.8 kcal mol 1 are found for the substitution and elimination, respectively. 

Studying the ESPT of hydroxy aromatic sulfonates, Huppert and co-workers [40-44] suggested an alternative model based on the geminate proton-anion recombination, governed by diffusive motion. The analysis was carried out by using Debye-Smoluchowskii-type diffusion equations. Their ESPT studies in water-methanol mixtures showed that solvent effects in the dissociation rate coefficient are equal to the effects in the dissociation equilibrium constant [45], 4-Hydroxy-1-naphthalenesulphonate in a water-propanol mixture as the solvent system has been found to behave somewhat differently than water-methanol or water-ethanol media, with a possible role of a water dimer [46,47],... 

The low equilibrium constant and the strongly nonideal behavior that causes the forming of the binary azeotropes methyl acetate/methanol and methyl acetate/ water make this reaction system interesting as a possible RD application (33). Therefore, methyl acetate synthesis has been chosen as a test system and investigated in a semibatch RD column. Since the process is carried out under atmospheric pressure, no side reactions in the liquid phase occur (146). 

In the half-cell of Eq. (5.24), the concentration of AgClj" must be small compared to that of Cl-, or a liquid-junction potential will result because the mobilities of AgClJ and Cl- are not the same. Thus, for a reference electrode of the second kind to be elfective in cells without appreciable junction potentials, the equilibrium constant for the reaction of Eq. (5.25) must be smaller than unity (preferably <0.1). In water, methanol, formamide, and V-methyl-formamide, this criterion is met, but in most organic solvents the equilibrium constant for the reaction of Eq. (5.25) ranges from 30 to 100. The silver chloride electrode is not recommended for general use in organic solvents.27... 

It is unfortunate that there has been so little work devoted to quantitative measurements of cation-pseudobase equilibria in methanol and ethanol since these media have several advantages over water for the determination of the relative susceptibilities of heterocyclic cations to pseudobase formation. The enhanced stability of the pseudobase relative to the cation in alcohols compared to water is discussed earlier this phenomenon will permit the quantitative measurement of pseudobase formation in methanol (and especially ethanol) for many heterocyclic cations for which the equilibrium lies too far in favor of the cation in aqueous solution to allow a direct measurement of the equilibrium constant. Furthermore, the deprotonation of hydroxide pseudobases (Section V,B) and the occurrence of subsequent irreversible reactions (Sections V,C and D), which complicate measurements for pKR+ > 14 in aqueous solutions, are not problems in alcohol solutions. Data are now available for the preparation of buffer solutions in methanol over a wide range of acidities.309-312 An appropriate basicity function scale will be required for more basic solutions. The series of -(substituted phenyl)pyridinium cations (163) studied by Kavalek et al.i2 should be suitable for use as indicators in at least some of the basic region. The Hm and Jm basicity functions313 should not be assumed90 to apply to methoxide ion addition to heterocyclic cations because of the differently charged species involved in the indicators used to construct these scales. 

A situation often encountered with thermodynamic data is that different sources may present different correlations for a given quantity. In this case study, for example, two different expressions are given for the equilibrium constant for the water-gas shift reaction. Equations 13.9 and 13.12. By what percentages do the conversions of CO and H2 and production of methanol differ if, instead of using Equation 13.12 in Problem 13.15(a), Equation 13.9 is used. Provide at least two reasons for the variations in the two equations for the water-gas shift equilibrium constant. 

Kooner, Z. S. Phutela, R. C. Fenby, D. V. Determination of the equilibrium constants in water-methanol deuterium exchange reaction from vapor pressure measurements. Aust. J. Chem. 1980, 33, 9-13. 

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