Hydration enthalpies

It will be noted that hydration enthalpy decreases with increasing ionic radius and increases very sharply with increase in ionic charge, these results being what we should expect for an electrostate interaction between a charged ion and the dipole of a water molecule (p, 44). 

The enthalpy of solution is quite small for many simple ionic compounds and can be either positive or negative. It is the difference between two large quantities, the sum of the hydration enthalpies and the lattice energy. 

Although the data for the silver halides suggest that silver(I) fluoride is likely to be more soluble than the other silver halides (which is in fact the case), the hydration enthalpies for the sodium halides almost exactly balance the lattice energies. What then is the driving force which makes these salts soluble, and which indeed must be responsible for the solution process where this is endothermic We have seen on p. 66 the relationship AG = — TAS and... 

Obviously sufficient energy is available to break the A1—Cl covalent bonds and to remove three electrons from the aluminium atom. Most of this energy comes from the very high hydration enthalpy of the AP (g) ion (p. 78). Indeed it is the very high hydration energy of the highly charged cation which is responsible for the reaction of other essentially covalent chlorides with water (for example. SnCl ). 

This is an exothermic process, due largely to the large hydration enthalpy of the proton. However, unlike the metallic elements, non-metallic elements do not usually form hydrated cations when their compounds dissolve in water the process of hydrolysis occurs instead. The reason is probably to be found in the difference in ionisation energies. Compare boron and aluminium in Group III ... 

The enthalpy changes AH involved in this equilibrium are (a) the heat of atomisation of the metal, (b) the ionisation energy of the metal and (c) the hydration enthalpy of the metal ion (Chapter 3). 

Heat of atomisation Sum of 1st and 2nd ionisation energies Hydration enthalpy AH... 

Bashin A A and K Namboodiri 1987. A Simple Method for the Calculation of Hydration Enthalpies c Polar Molecules with Arbitrary Shapes. Journal of Physical Chemistry 91 6003-6012. 

Table 1.3 Esti mated values of the four components of the contribution made by ligand field stabilization energy to the lattice enthalpy of KsCuFe, to the hydration enthalpy of Ni (aq), AH (Ni, g), and to the standard enthalpy change of reaction 13.

In KCl, the time dependence showed a marked decrease as the concentration was increased, becoming practically independent of time at a concentration of 1.0 M. This could be explained by adsorbed layers of cations on the surfaces, which reduce the attractive force between them. Further, this reduction with concentration became greater as the hydration enthalpy of the cations increased. This was interpreted in terms of the observations of Hi-... 

FIG. 19 Proposed mechanism for the difference of adhesive force between cations of low and high hydration enthalpies. (Reprinted from Ref. 88. Copyright 2000 by Academic Press.)... 

Figure 13. Hydration enthalpy, AHo,, for hydration of H3N(CH2)pNHf+ plotted versus 1/p+2 which is approximately proportional to the reciprocal of the distance between the two charged centers. This leads to a straight line. Such a relationship can be expected because the Coulombic energy is proportional to the reciprocal of the distance.

The hydration enthalpy of the Al3+ ion is enormous (-4690k) mol-1), and there are some interesting effects produced as a result. When NaCl is dissolved in water and the solvent evaporated, the solid NaCl can be recovered. If A1C13 is dissolved in water, evaporation of the water does not yield the solid A1C13. The Al3+ ion is so strongly solvated that other reactions become energetically more favorable than removing the solvent. This can be shown as follows. 

Cations in aqueous solutions have an effective radius that is approximately 75 pm larger than the crystallographic radii. The value of 75 pm is approximately the radius of a water molecule. It can be shown that the heat of hydration of cations should be a linear function of Z /r where is the effective ionic radius and Z is the charge on the ion. Using the ionic radii shown in Table 7.4 and hydration enthalpies shown in Table 7.7, test the validity of this relationship. 

Chloride hydrate Enthalpy (kJ mol-1) Chloride hydrate Enthalpy (kJ mol-1)... 

Figure 2.5 The Born-Flaber cycle for lithium methoxide and the hydration enthalpies of the ions.

Reactive ionic compounds are therefore useless to derive hydration enthalpies (or more generally, solvation enthalpies). Fortunately, there are many alternatives. Take lithium chloride, for example, and data from the NBS Tables [ 17]. The enthalpy of solution of this solid in water, at infinite dilution, is given by... 

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