Table 3.4 Some typical buffer solutions
Solutions pH Range
phthalic acid and potassium hydrogen phthalate 2.2–4.2
citric acid and sodium citrate 2.5–7.0
acetic acid and sodium acetate 3.8–5.8
sodium dihydrogen phosphate and disodium hydrogen phosphate 6.2–8.2
ammonia and ammonium chloride 8.2–10.2
borax and sodium hydroxide 9.2–11.2solution. To facilitate this practical approach, equation (3.20) may be written in a different form,
The approximate pH range over which a buffer solution remains effective can be deduced from Figure
3.1. The limits of effective buffering can be seen as the points at which the ratio [AH]/[A–] becomes
10:1 or 1:10 whence substitution in equation (3.20) yields
The pH of Salt Solutions
Table 3.4 shows some typical buffer solutions.
When an acid in solution is exactly neutralized with a base the resulting solution corresponds to a
solution of the salt of the acid-base pair. This is a situation which frequently arises in analytical
procedures and the calculation of the exact pH of such a solution may be of considerable importance.
The neutralization point or end point in an acid-base titration is a particular example (Chapter 5). Salts
may in all cases be regarded as strong electrolytes so that a salt AB derived from acid AH and base B
will dissociate completely in solution. If the acid and base are strong, no further reaction is likely and
the solution pH remains unaffected by the salt. However if either or both acid and base are weak a more
complex situation will develop. It is convenient to consider three separate cases, (a) weak acid-strong
base, (b) strong acid-weak base and (c) weak acid-weak base.
(a)—
Weak Acid-strong Base Solutions
The conjugate base A– will react with water and undergo hydrolysis,
producing undissociated acid and hydroxyl ions with an accompanying rise in pH. The equilibrium
constant for this reaction is known as the hydrolysis constant Kh.