Enthalpy of exchange

TABLE V. Calorimetrically measured enthalpies of exchange of pyridinium and bipyridium salts in... 

Differential enthalpies of exchange, plotted as a function of K saturation, exhibit a stepped character (Figure 1 and Table II). This can only reflect different types of adsorption sites on the clay surfaces, the group of sites with the most negative enthalpy... 

From Equation 3 we calculate the free enthalpy of exchange AG2CiPo(V-------RT lna. 

The relationship between activation energies and the expected effects of substituents has also been examined (63JA329). Lack of correlation indicates reaction on the conjugate acid, so that the activation energy includes the enthalpy of dissolution and the activation enthalpy of exchange. Appropriate correction for the former then leads to satisfactory Hammett plots. 

TABLE 4.2. Thermodynamic Equilibrium Constants of Exchange (ATCJ and Standard Enthalpy of Exchange (AH°e ) Values for Binary Exchange Processes on... 

Also, from Keq, the standard enthalpy of exchange, AH° can be calculated (Table 4.2) using the van t Hoff equation ... 

K represents the real equilibrium constant for an exchange between Z = 0 and Z = Zmax. The standard free energy of exchange is obtained as AG = — RT In K. The enthalpy of exchange is derived from the Vant Hoff equation, and the entropy change is calculated from the AG and AH values ... 

The deterniination of the enthalpy of exchange can be obtained directly by micro-calorimetry using two methods [14] ... 

Enthalpy of exchange between monovalent counterions and the cation TEA as reference... 

Fig. 5. Enthalpy of exchange TEA /Li obtained by microcalorimetry as a function of the charge density.

If > x, the excess isotherm is practically linear between Xi = 0.1 and 0.7, and the enthalpies of exchange measured are moderately endothermic. If the molar fraction of the bulk phase is larger than Xi = 0.8, the displacement of benzene by methanol wiU preferentially occur, resulting in a considerable endothermic effect on the porous apolar surface. 

Using the same technique, MacDonald and Evans (2002) measured the enthalpy of exchange of phenol-water from dilute aqueous solutions on commercial BPL carbons with different oxygen contents. They found that a strong preference of water to H-bond to surface oxygen complexes lowers the phenol exchange enthalpy from aqueous solution because the phenol exchange enthalpy was inversely proportional to the net molar enthalpy of immersion into water. 

---