Other common examples are aluminium chloride, stannic chloride,
zinc chloride, etc. We shall, at this point, be concerned essentially
with proton acids and the effect of structure on the strength of a
number of organic acids and bases will now be considered in turn.
Compounds in which a C—H bond is ionised will be considered
subsequently (p. 210), however.
ACIDS
(0 pKaThe strengtb^of an acid, HX, in water, i.e. the extent to which it is
dissociated, may be estimated oy considering the equilibrium:
HaO:+HX ^ HsO® + X®
Then the equilibrium constant is jiven by
fKa~ [HX]
jjje concentration of water being taken as constant as it is present in
such large excess. It should be emphasised that Ka, the acidity constant
of the acid in water, is only approximate if concentrations instead of
activities have been used. The constant is influenced by the composi
tion of the solvent in which the acid is dissolved fjsee below) and by
other factors but it does, nevertheless, serve as a useful-guide t<^cid
strength. In order to avoid writing negative powers of 10, Ka is
generally converted into $Ka (j>Ka = -log 10 KJ; thus while Ka for
acetic acid in water at 25° is 1 • 79 x 10~^5 , pK 0 = 4 • 76. The smaller the
numerical value of pK 0 , the stronger is the acid to which it refers.(ii) Effect of solvent
The influence of the solvent on the dissociation of acids (and of bases)
can be profound; thus hydrogen chloride which is a strong acid in
water is not ionised in benzene. Water is a most effective ionising
solvent on account (a) of its high dielectric constant, and (b) of its
ion-solvating ability. The higher the dielectric constant of a solvent the
smaller the electrostatic energy of any ions present in it; hence the
more stable such ions are in solution.
Ions in solution strongly polarise solvent molecules near them and
the greater the extent to which this can take place, the greater the
stability of the ion, which is in effect stabilising itself by spreading
its charge. Water is extremely readily polarised and ions stabilise