Page 60 - Vogel's TEXTBOOK OF QUANTITATIVE CHEMICAL ANALYSIS

2 FUNDAMENTAL THEORETICAL PRINCIPLES OF REACTIONS IN SOLUTION
the properties of their solutions, and deviations are apparent from ideal laws
(which are assumed in the derivation of the Mass-Action Law by thermodynamic
or kinetic methods); the deviations from the ideal laws are usually expressed
in terms of activities or activity coefficients. For Our purpose, the deviations due
to interionic attractions and ionic activities will be regarded as small for small
ionic concentrations and the equations will be regarded as holding in the same
form at higher concentrations, provided that the total ionic concentration does
not Vary much in a given set of experiments.
To use the above expression for measuring the strength of an acid, a standard
acid-base pair, Say A,-B,, must be chosen, and it is usually convenient to
refer acid-base strength to the solvent. In water the acid-base pair H30+-H20
is taken as the standard. The equilibrium defining acids is therefore:

and the constant

gives the strength of A, that of the ion H30+ being taken as unity. Equation (c)
represents what is usually described as the dissociation of the acid A in water,
and the constant K' is closely related to the dissociation constant of A in water
as usually defined and differing only in the inclusion of the term [H ,O] in the
denominator. The latter term represents the 'concentration' of water molecules
in liquid water (55.5 moles per litre on the ordinary volume concentration scale).
When dealing with dilute solutions, the value of [H,O] may be regarded as
constant, and equation (6) may be expressed as:

by writing H + for H30+ and remembering that the hydrated proton is meant.
This equation defines the strength of the acid A. If A is an uncharged molecule
(e.g. a weak organic acid), B is the anion derived from it by the loss of a proton,
and (7) is the usual expression for the ionisation constant. If A is an anion such
as H, PO;, the dissociation constant [HPO:-] [H + ]/CH, PO; ] is usually
referred to as the second dissociation constant of phosphoric(V) acid. If A is a
cation acid, for example the ammonium ion, which interacts with water as shown
by the equation

the acid strength is given by [NH,][H+]/[NH,+].
On the above basis it is, in principle, unnecessary to treat the strength of
bases separately from acids, since any protolytic reaction involving an acid must
also involve its conjugate base. The basic properties of ammonia and various
amines in water are readily understood on the Br~nsted-Lowry concept.