We often think of ionic compounds as being water soluble. This is a simplification: most ionic compounds are only sparingly soluble. As such, their solubility can be treated as an equilibrium reaction and all the tools we have learned can be applied.
First, a review from CHM 101: we need to know the solubility rules.
You must memorize these rules because they are the guide to predicting precipitation reactions.
Precipitation reactions occur when two ions are present in a solution that can combine to form a sparingly soluble salt.
Complete and balance the following reactions:
Pb(NO3)2(aq) + K2SO4(aq) ?
Fe(NO3)3(aq) + CsCl(aq) ?
Co(NO3)2(aq) + H2S(aq) ?
The solubilization of a sparingly soluble salt, MpXq, can be written as an equilibrium reaction:
MpXq(s) → ← pMq+(aq) + qXp–(aq)
The equilibrium constant is
Kc = [Mq+]ep[Xp–]eq = Ksp
where Ksp is the solubility product constant.
Just as with Ka or Kb, the subscript defines the reaction.
Solubility equilibria and their associated equilibrium constants are nearly always approximations. This is because we often do not know the actual ions in solution. Despite this, we can use Ksp values to obtain reasonable estimates of the solubility of ionic salts. Values for Ksp are tabulated in the Table of Solubility Product Constants.
Compare the molar solubilities of barium sulfate, lead(II) chloride, and magnesium arsenate.
Strategy: For each salt, write the solubilization reaction, write the mass action expression, set up a table of concentrations, plug the equilibrium expressions into the mass action expression, solve. (A typical equilibrium problem.)
Compare:
Compound Ksp Solubility (M)
BaSO4 1.1×10–10 1.0×10–5
PbCl2 1.6×10–5 1.6×10–2
Mg3(AsO4)2 2×10–20 5×10–5
Note that magnesium arsenate is more soluble than barium sulfate even though the equilibrium constant for barium sulfate is ten billion times smaller than the equilibrium constant for magnesium arsenate.
Warning: you cannot use Ksp directly to compare solubilities.