base dissociation constant
/ BAYS dis-oh-see-AY-shun KON-stunt /
If the acid dissociation constant scores how eagerly an acid gives away a proton, the base dissociation constant scores the mirror image: how eagerly a base grabs one. Picture an open hand reaching out — some bases snatch a hydrogen ion readily, others barely manage to catch one.
Formally, for a base B in water, the base dissociation constant Kb is the equilibrium constant for the reaction in which B takes a hydrogen ion from water to form its conjugate acid plus a hydroxide ion. It equals the product of the conjugate acid and hydroxide concentrations divided by the concentration of the unreacted base. A large Kb marks a strong, avid base; a small one a weak, hesitant base, and it is often quoted as pKb.
This constant matters because it predicts how basic a solution will become and where a weak-base titration will turn. The honest and very useful caveat is that you rarely need to look Kb up separately: for any conjugate acid-base pair, Ka times Kb equals the ion product of water Kw, so pKa plus pKb equals about 14, and knowing one immediately hands you the other.
Ammonia has a Kb of about 1.8 times ten to the minus five, so it is a weak base; its conjugate acid, the ammonium ion, accordingly has a pKa near 9.25, and the two pK values add up to about 14.
For a conjugate pair, pKa plus pKb adds up to about 14 in water.
Because Ka times Kb equals Kw, a stronger base always has a weaker conjugate acid and vice versa. Modern tables increasingly list only Ka or pKa for the conjugate acid, expecting you to convert, so it pays to be fluent with this relation.