is incorporated in the constants Ka and Kb. If the concentrations are expressed in mol dm–^3 , Ka and Kb
have the same units. Table 3.3 lists pKa values for some typical compounds. The behaviour of the
conjugate base may be represented in line with equation (3.15), i.e.
Kw is the self-ionization constant for water (Table 3.2) and equation (3.18) reflects the not surprising
inverse relation between Ka and Kb. It is only when Ka and Kb for a compound are of different
magnitudes that it may be classified as an acid or a base. An example which is difficult to classify is
hypoiodous acid (HOI) where Ka = 2.5 × 10 –^11 mol dm–^3 and Kb = 3.2 × 10 –^10 mol dm–^3. Although Kb has
been widely used in the past, it is a quantity which is largely redundant, for Ka (or pKa) may be used to
express the strength of bases as well as acids, see Table 3.3.
Table 3.3 shows that values of Ka and Kb vary over a wide range and also that there is no clear dividing
line between strong acids and weak acids or strong bases and weak bases. However as a rough guide
weak acids may be regarded as those having values of pKa in the range 4–10, those having pKa = 4
being called 'strong weak acids' and those with pKa = 8–10 'very weak acids'. The pH of a solution of a
strong acid or base may be related directly to the concentration of the acid or base. However, weak acid
or base systems present a rather more complex pattern.
Weak Acid and Weak Base Equilibria
Equation (3.14) may be rewritten
and presented in a log form using the 'p' notation
This is a useful equation as it gives the relation between the pH of a solution, the dissociation constant
for the acid, and the composition of the solution. When the relation is represented graphically (Figure
3.1) some further valuable points emerge. As it is the difference between the pH of the solution and pKa
for the acid which is important rather than the pH