Introductory University Chemistry I

Solutions and Solubility

Solubility of Substances in Water

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. Almost all sodium, potassium, and ammonium salts are highly soluble; the only significant exception is KClO₄, which is moderately soluble but is still sometimes used to precipitate potassium ion from aqueous solutions. Metal nitrates are soluble almost without exception. Metal carbonates and phosphates are generally insoluble or slightly soluble, with the exception of those of sodium, potassium, and ammonium which are highly soluble; magnesium ammonium phosphate is used for the precipitation of magnesium ion.

Metal halides are generally highly soluble, with the exception of those of silver, lead, and mercury (I). Lead chloride is slightly soluble while silver and mercury (I) chlorides are much less soluble. Sulfate salts are generally highly soluble as well, with more exceptions; calcium, barium, strontium, lead, and mercury (I) sulfates are almost insoluble while silver sulfate is slightly soluble. Metal sulfides are generally insoluble in water.

Solid-Solution (Solubility) Reactions

When solids dissolve, the solutes are no longer pure substances and their 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 and are taken up under the following main heading. 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-) = K_sp. If the molar concentrations are taken as good approximations to the activities, which in dilute solutions they are, then K_sp = [Ag+][Cl-].

Example. Let us write and simplify to the extent possible the equilibrium constant for the equilibrium Al³⁺(aq) + 3OH^-(aq) <--> Al(OH)₃(s) For this equilibrium K = 1/[Al³⁺] [OH^-]³ = 1/K_sp. where K has the units dm¹²/mol⁴, or (dm³)⁴/mol⁴.

The form of equilibrium constant indicated as K_sp 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.

Figure is not available.

The solubility of solid salts in water, and in most other solvents, increases with temperature while that 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:

NaCl(c) <--> Na⁺(aq) + Cl^-(aq); H0 = (-240.12 - 167.159) - (-411.153) = +3.87 kJ/mol

KNO₃(c) <--> K⁺(aq) + NO₃^-(aq); H0 = (-252.38 - 205.0)-(-494.63) = +37.3 kJ/mol

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