Dissolution of Ionic Solids in Water

Solvation energies of ions with unit charges are of the order of -100 kcal/mol, and therefore, the combined solvation energies can lower the energy of solid ionic materials by ca. 200 kcal/mol, which usually exceeds the Coulomb attractions in the solid (which are lost during dissolution) for monocharged ions. In such a case, the solvation of the solid will be exothermic, and the excess energy be released as heat. 

The Resistance of Ionic Materials with Multiply Charged Ions to Dissolve  

Ca C03 does not dissolve in water. The reason is rather simple and can be seen if we use the equation Eqiqi = 332 [Qi x Q Ri2 for the attraction of the ions. Even if we assume that the distance R 2 is identical to the one in Na+Cl (it is shorter in fact), the attraction energy for Ca +C03 will be huge, -476 kcal/mol, and in the solid (assuming the same structure as in Na Cl and hence also the same M = 1.75 value), this quantity will become -833 kcal/mol. Now the solvation energy of the ions will not compensate for the loss of the attraction in the solid, and so the solid will resist dissolution. 

when no energetic effect (of breaking bonds and making new ones, as in the dissolution of an ionic material in water) is involved, the most probable situation we anticipate to find is that which has maximum entropy/disorder. 

Getting back to the dissolution of an ionic solid in water, it is clear that going from the nicely ordered ionic solid to separate ions in solution, arranged in many different and random patterns, increases the entropy/disorder, and hence, solvation will be preferred unless the attraction in the solid is much stronger than the attractions felt by each ion from the water molecules that surround it. When the interaction energies in the solid and the solution are not too different, then endothermic dissolution can nevertheless happen, and it will be driven by the increase of entropy/disorder. This is the case of the dissolution of [NH4+Cl ]soiid in water. 

---

To understand the dissolution of ionic solids in water, lattice energies must be considered. The lattice enthalpy, A Hh of a crystalline ionic solid is defined as the energy released when one mole of solid is formed from its constituent ions in the gas phase. The hydration enthalpy, A Hh, of an ion is the energy released when one mole of the gas phase ion is dissolved in water. Comparison of the two values allows one to determine the enthalpy of solution, AHs, and whether an ionic solid will dissolve endothermically or exothermically. Figure 1.4 shows a comparison of AH and A//h, demonstrating that AgF dissolves exothermically. 

---