Page 59 - Vogel's TEXTBOOK OF QUANTITATIVE CHEMICAL ANALYSIS

STRENGTHS OF AClDS AND BASES 2.13

2.12 ACID-BASE EOUlllBRlA IN WATER
Consider the dissociation of a weak electrolyte, such as acetic acid, in dilute
aqueous solution:
CH3COOH + H20 e H30+ +CH,COO-
This will be written for simplicity in the conventional manner:
CH,COOH + H+ + CH,COO-
where H+ represents the hydrated hydrogen ion. Applying the Law of Mass
Action, we have:
[CH,COO-] x [H+]/[CH,COOH] = K
K is the equilibrium constant at a particular temperature and is usually known
as the ionisation constant or dissociation constant. If 1 mole of the electrolyte
is dissolved in Vlitres of solution (V = l/c, where c is the concentration in
moles per litre), and if cc is the degree of ionisation at equilibrium, then the
amount of un-ionised electrolyte will be ( 1 - cc) moles, and the amount of each
of the ions will be cc moles. The concentration of un-ionised acetic acid wiil
therefore be ( 1 - cc)/ and the concentration of each of the ions cc/ K Substituting
in the equilibrium equation, we obtain the expression:
cc2/(1-cc)V = K or cc2c/(l-cc) = K
This is known as Ostwald's Dilution Law.
Interionic effects are, however, not negligible even for weak acids and the
activity coefficient product must be introduced into the expression for the
ionisation constant:

Reference must be made to textbooks of physical chemistry (see Bibliography,
Section 3.39) for details of the methods used to evaluate true dissociation
constants of acids.
From the point of view of quantitative analysis, sufficiently accurate values
for the ionisation constants of weak monoprotic acids may be obtained by using
the classical Ostwald Dilution Law expression: the resulting 'constant' is
sometimes called the 'concentration dissociation constant'.

2.13 STRENGTHS OF AClDS AND BASES
The Br~nsted-Lowry expression for acid-base equilibria (see Section 2.4)

A, +B2=A2+Bl (b)
leads, upon application of the Law of Mass Action, to the expression:

where the constant K depends on the temperature and the nature of the solvent.
This expression is strictly valid only for extremely dilute solutions: when ions
are present the electrostatic forces between them have appreciable effects on