James Richard Fromm
Water is an excellent solvent for many compounds. Some dissolve in it as molecules while others, called electrolytes, dissociate and dissolve not as neutral molecules but as charged species called ions. Compounds which exist as solid ionic crystals dissolve in water as ions, and most of them are highly soluble in water. "Highly soluble" is a somewhat elastic description, but generally means soluble to at least the extent of forming 0.1 to 1.0 molar aqueous solutions. Salts which are less soluble in water than this at room temperature are called slightly soluble salts.
The solubility of an ionic salt depends both upon its cations and its anions, but for simple salts in aqueous solution at room temperature the following general observations are useful.
When solids dissolve, the solutes are no longer pure substances and their chemical activity can no longer be taken as unity. In dilute solutions, aqueous or otherwise, activities of solutes are often taken as equal to their molar concentrations. These equilibria are called solubility equilibria. The example below shows how the form in which they are written compares to other equilibrium constants.
Example. The equilibrium constant for the reaction
AgCl(s) Ag+(aq) + Cl-(aq)
is written as
K = a(Ag+)a(Cl-)/a(AgCl);
more commonly, it is written in the form
Ka(AgCl) = a(Ag+)a(Cl-) = Ksp.
If the molar concentrations are taken as good approximations to the activities, which in dilute solutions they are, then Ksp = [Ag+][Cl-].
Example. Let us write and simplify to the extent possible the equilibrium constant for the solubility equilibrium
Al3+(aq) + 3OH-(aq) Al(OH)3(s).
For this equilibrium K = 1/[Al3+][OH-]3 = 1/Ksp. where Ksp has the units dm12/mol4, or (dm3)4/mol4.
The form of equilibrium constant indicated as Ksp is called the solubility product constant or, more commonly, the solubility product. This constant therefore must refer to the process of a solid going into solution (solubility) rather than the reverse, precipitation of solid from solution. As a consequence, the ions are products and appear in the numerator.
The value of the solubility product is temperature-dependent and is generally found to increase with increasing temperature. As a consequence, the molar solubility of ionic salts generally increases with increasing temperature. The extent of this increase varies from one salt to another. It is sometimes possible to take advantage of the difference in the effect of temperature to separate mixtures of different soluble salts. As the chart in the following Figure shows, a solution originally of equal concentration in KClO3 and KNO3 should upon heating and evaporation of water precipitate
KClO3 because KNO3 is by far the more soluble near the boiling point of water. The solubility of solid salts in water, and in most other solvents, increases with temperature while the solubility of gases decreases. This is an application of Le Chatelier's Principle, discussed in a preceding section. The heat or enthalpy change of the dissolution reaction for most solids is positive so the dissolution reaction is endothermic. For some solids, such as NaCl, the heat of solution is very small and so the effect of temperature is small also. For other salts, such as KNO3, the effect of temperature is much larger. This is so because the enthalpy change for dissolution of KNO3 is almost ten times as large as is the enthalpy change for dissolution of NaCl:
NaCl(c) Na+(aq) + Cl-(aq); DH0 = (-240.12 - 167.159) - (-411.153) = +3.87 kJ/mol
KNO3(c) K+(aq) + NO3-(aq); DH0 = (-252.38 - 205.0) - (-494.63) = +37.3 kJ/mol