Hydration, ionic, enthalpy

More recently, ionic enthalpies and free energies of hydration have been determined relative to jj+.88 168 171 This can be done, for example, by means of appropriate thermodynamic cycles.170,171 Then by assigning the AHhydTation and AGhydration of H+ on some basis, these quantities can be evaluated for the other ion. Rashin and Honig took AHhydTation (H+) to be -262.18 kcal/mole,88 while Florian and Warshel assumed that AGhydratj0n (H+) is -259.5 kcal/mole.171... 

Table 9 compares ionic enthalpies of hydration from the Bernal and Fowler,164 Latimer et al.165 and Rashin and Honig88 procedures. Given the inherent uncertainty, the latter two sets of data are remarkably similar, considering that they were obtained 46 years apart. A number of tabulations of the thermodynamic solvation properties of ions in various solvents have now appeared. It is important to keep in mind, however, that there is a degree of arbitrariness associated with the experimental AHsoivation and AGSoiVation of individual ions. 

Experimental Ionic Enthalpies of Hydration Obtained By Three Different Methods, in kcal/mole ... 

In Table 10 are listed some computed ionic enthalpies and free energies of hydration, as well as the corresponding experimental ones. The range in some of the latter illustrates the uncertainty associated with the experimental data, discussed in the previous Section. This must of course be taken into account in evaluating the calculated results. 

The ionic hydration energies (enthalpies and Gibbs free energies) of metals are consequently roughly a linear function of the square of the oxidation state divided by the effective ion radius (z /r ff. Figure 8.9).It may be added that AH y and AG y, respectively, of the individual alkali halide ion pairs form two linear branches with a maximum at Csl (KBr). 

The stable form of an ionic solid at saturation is often a hydrate. The enthalpy of solution for hydrates is generally positive (endothermic), especially at... 

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. 

Prediction of solubility for simple ionic compounds is difficult since we need to know not only values of hydration and lattice enthalpies but also entropy changes on solution before any informed prediction can be given. Even then kinetic factors must be considered. 

The solubilities of the ionic halides are determined by a variety of factors, especially the lattice enthalpy and enthalpy of hydration. There is a delicate balance between the two factors, with the lattice enthalpy usually being the determining one. Lattice enthalpies decrease from chloride to iodide, so water molecules can more readily separate the ions in the latter. Less ionic halides, such as the silver halides, generally have a much lower solubility, and the trend in solubility is the reverse of the more ionic halides. For the less ionic halides, the covalent character of the bond allows the ion pairs to persist in water. The ions are not easily hydrated, making them less soluble. The polarizability of the halide ions and the covalency of their bonding increases down the group. 

Morris DFC (1968/1969) An Appendix to Structure and Bonding. 4 6 157-159 Morris DFC (1968) Ionic Radii and Enthalpies of Hydration of Ions. 4 63-82 Mortensen OS (1987) A Noncommuting-Generator Approach to Molecular Symmetry. 68 1-28... 

If a substance is to be dissolved, its ions or molecules must first move apart and then force their way between the solvent molecules which interact with the solute particles. If an ionic crystal is dissolved, electrostatic interaction forces must be overcome between the ions. The higher the dielectric constant of the solvent, the more effective this process is. The solvent-solute interaction is termed ion solvation (ion hydration in aqueous solutions). The importance of this phenomenon follows from comparison of the energy changes accompanying solvation of ions and uncharged molecules for monovalent ions, the enthalpy of hydration is about 400 kJ mol-1, and equals about 12 kJ mol-1 for simple non-polar species such as argon or methane. 

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. 

It is quite difficult to measure an accurate enthalpy of solution A//( olutioni with a calorimeter, but we can measure it indirectly. Consider the example of sodium chloride, NaCl. The ions in solid NaCl are held together in a tight array by strong ionic bonds. While dissolving in water, the ionic bonds holding the constituent ions of Na+ and Cl- in place break, and new bonds form between the ions and molecules of water to yield hydrated species. Most simple ions are surrounded with six water molecules, like the [Na(H20)6]+ ion (VI). Exceptions include the proton with four water molecules (see p. 235) and lanthanide ions with eight. 

Energy is needed to break the ionic bonds in the solid salt and energy is liberated forming hydration complexes like VI. We also break some of the natural hydrogen bonds in the water. The overall change in enthalpy is termed the enthalpy of solution, A// olutioni. Typical values are —207 kJmol-1 for nitric acid 34 kJmol-1 for potassium nitrate and —65.5 kJmol-1 for silver chloride. 

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