Understandings:
• The expression for the dissociation constant of a weak base (Kb).
• For a conjugate acid base pair, Ka × Kb = Kw.
• The relationship between Kb and pKb is (pKb = -log Kb).
Applications and skills:
• Solution of problems involving [H+ (aq)], [OH–(aq)], pH, pOH, Ka, pKa, Kb and pKb.
• Discussion of the relative strengths of acids and bases using values of Ka, pKa, Kb and pKb.
Guidance:
• Only examples involving the transfer of one proton will be assessed.
• Students should state when approximations are used in equilibrium calculations.
When bases interact with water they do so by removing a hydrogen ion creating hydroxide ions in solution.
NH3 (aq) + H2O ⇋ NH4+(aq) + OH-(aq)
Once again, the water concentration is effectively constant and allows us to define a new constant, Kb:
Kb = [NH4+] [OH-]
[NH3]
Kb is referred to as the base equilibrium constant and gives a measure of the extent of the equilibrium. Large values for Kb means strong base. However, as for acids, the values are usually very small and the -log(10) form of the information is used to generate convenient sized numbers.
pKb = - log(10) Kb
The lower the pKb value the stronger the base.
With bases it may be more convenient to use the conjugate acid of the base to define a new Ka value, this time for the reverse reaction. In the ammonia base equilibrium above, the conjugate acid is the ammonium ion, NH4+, whose dissociation proceeds according to the following equation:
NH4+ + H2O ⇋ NH3 + H3O+
For which the acid dissociation constant is:
Ka = [NH3] [H+]
[NH4+]
This allows us to have a pKa value for a specified base, although it is actually the pKa value of the conjugate acid of the base.
Ka x Kb = Kw
This means (at Room Temperature) Ka x Kb = 1x10-14
And if we know this then that means we ca use all of the following relationships to calculate anything we need about acids and bases: