Bronsted-Lowry acids- proton donors, they release hydrogen ions (H+) when they're mixed with water.
HA (aq) <=> H+ (aq)+ A- (aq)
H+ ions are never by themselves in water they're always combined with H2O to form hydroxonium ions, H3O+.
HA (aq) +H2O (l) <=> H3O+ + A- (aq)
Bronsted-Lowry bases- proton acceptors.
When they're in solution, they accept hydrogen ions from water molecules and form (OH-) ions.
B (aq) + H2O (l) <=> BH+(aq) + OH- (aq)
Lewis acid and bases
Lewis Acid - lone pair acceptor e.g. BF3, H+, ALCl3
Lewis Base - lone pair donor NH3, H20
Strong/weak acids and bases
Strong acids/bases- dissociate (or ionise) completely into ions in aqueous solution.
These reactions are actually reversible reactions but the equilibrium lies very far to the right, so only the forward reaction is shown in the equation.
Weak acids- partially dissociate into ions in aqueous solution, so only small numbers of H+ ions are formed.
Weak bases-partially dissociate into ions in aqueous solution.
An equilibrium is set up which lies to the left.
Monoprotic, Diprotic and Triprotic acids
Monoprotic - release one proton. E.g. HCl, HNO3
Diprotic - release two protons. E.g. H2SO4
Triprotic- release three protons. E.g. H3PO4
Conjugate pairs
Conjugate acid/base pairs differ by H+.
The species that has lost a proton is the conjugate base.
The species that has gained a proton is the conjugate acid.
HA + B <=> BH+ +A-
When HA loses a proton it forms A- and when A- gains a proton it forms HA.
HA and A- are linked by proton transfer so are a conjugate pair.
HA has gained a proton so is the conjugate acid of A.
A has lost a proton so is the conjugate base of HA.
The base, B, takes a proton from the acid, HA, to form BH+
B is the conjugate base of BH+ and
BH+ is the conjugate acid of B.
When an Acid and a Base react you get 2 conjugate pairs:
Water is amphoteric- it behaves as both an acid and a base.
