Solubility
Equilibrium
Using the solubility rules one
would predict that if a solution of sodium hydroxide and magnesium chloride
were mixed, a precipitate of magnesium hydroxide would form. But, strictly
speaking, it should be stated that a precipitate may form. If the
solutions mixed were very dilute a precipitate might not form. So, what
concentrations of ions will give a precipitate? We can answer this question by
considering the equilibrium between a solid salt and its ions in a saturated
solution of the salt (solubility equilibrium) in a more quantitative manner.
Solubility
Equilibrium: The equilibrium between the solid and its ions. Example, when calcium sulphate is
placed in water a small amount dissolves and we obtain the equilibrium:
CaSO4(s) ---à Ca2+(aq) + SO42-(aq)
The expression for the equilibrium constant is:
Kc = [Ca2+(aq)][SO42-(aq)]
[CaSO4(s)]
However, this is an example of a heterogeneous equilibrium. Therefore,
the concentration of the pure solid, CaSO4, is a constant and can be
incorporated into the equilibrium constant to give another constant, Ksp,
which is called the solubility product constant.
Solubility
Product Constant, Ksp: The equilibrium constant, Ksp, for the
solubility of a salt that, for a saturated solution, is equal to the product of
the concentrations of the ions involved in the equilibrium, each raised to the
power of its coefficient in the equation for the equilibrium.
Example: Write the solubility product constant expression for Ag2CO3.
Calculating
The Solubility Product Constant, Ksp, From Solubility Data
The value of the solubility
product constant, Ksp, for an ionic compound can be obtained by
measuring the salts molar solubility in water, the number of moles of solute
needed to dissolve, to give one litre of a saturated solution.
Example: One litre of water
will dissolve 4.9 x 10-3 mol of CaSO4. What is the
Ksp for CaSO4?
Note: Very insoluble salts have very small Ksp
values
More soluble salts have larger Ksp values