A strong acid, with pKa<0, is almost completely ionized in aqueous solution. Similarly, a strong base is almost fully protonated in solution.

A weak acid, with pKa>0, is not completely ionized in solution, and a weak base is only partially protonated.

Note that it is important to distiguish between a weak base which is only partially protonated in solution, and the conjugate base of a very strong acid, which would such a weak base that it is negligibly protonated in solution, and essentially shows no basic properties at all.

Strong Acid pKa<0
Strong Base pKa>14
Weak Acid pKa=4.75
Weak Acid pKa=1.92
Weak Base pKa=9.25

When an acid reacts with a base, the pH of the equivalence point, or stoichiometric point, where equal volumes of the acid and base have been reacted, can be predicted by the acid strengths, pKa‘s, of the two reactants.

If a strong acid (HCl) reacts with a strong base (NaOH), the only ions present are the cation from the base (Na+), and the anion from the acid (Cl), which have little influence on the pH, and the H3O+ and OH ions from the autoprotolysis of water, and as these are equal the pH of equivalence is equal to 7.

When a weak acid reacts with a strong base, at equivalence there will be the countercation from the base (Na+), and the counter anion from the acid (CH3COO). However, the conjugate base of weak acid will be a weak base, and its presence will mean that a pH greater than 7 will be observed, giving a basic solution.

In the above example, the ethanoate ion will become partially protonated from the water (an equilibrium is set up), and so the solution contains a slight excess of hydroxide ions.

When a strong acid reacts with a weak base, at equivalence there will be the counter cation from the base (NH4+), and the counter anion from the acid (Cl). However, the conjugate acid of weak base with be a weak acid, and its presence will mean that a pH less than 7 will be observed, giving an acidic solution. In this example, the ammonium ion gives up some protons to the solvent, meaning that there is a small excess of H3O+ in solution.

When an weak acid-strong base titration is performed, the pH will vary with the amount of base added to the acid. The pH can be determined from the acid strength, pKa, and the concentrations of the acid and base, using the Henderson-Hasselbach equation.

When the molar concentrations of the acid and base are equal, the pH of the solution is equal to pKa. Hence, at the equivalence point, the pH is given by the pKa of the acid.

If the titration is continued, and excess base is added, the pH of the solution can be given in terms of the concentration of excess base, [B*].

The variation of the pH with amount of base added can be expressed in terms of a pH curve. The example below shows a pH curve for the titration of a weak acid with a strong base.

A0 = initial concentration of acid

A’ = concentration of acid

S = concentration of salt (product of reaction)

B’ = concentration of excess base

Similar curves can be drawn for the other strong-weak, acid-base combinations, reflecting the changed value of the pH at the stoichiometric point, but they all have similar shapes.

Buffer Solutions

When the concentrations of acid and salt are similar, the pH of the solution varies only slowly with the amount of base added (the horizontal region in the above pH curve, when pH = pKa). This is the basis of a buffer solution: a solution which can oppose changes in pH when small amounts of strong acids and strong bases are added to the solution.

In the buffer solution, where there is a lot of salt present, there is an excess of the conjugate acid A, and hence when a strong acid is added, HA can be produced which reduces the H3O+ concentration, and hence restores the pH towards the initial value. Furthermore, when strong base is added, the HA molecules which are present can provide H3O+ to react with the base and remove it, and hence restore the pH downwards.

The action of a buffer solution, in opposing the change in pH it is subjected to, is an example of Le Chatelier’s principle.

This states that a system at equilibrium, when subjected to a disturbance, responds in a way that tends to minimize the effect of the disturbance.