An acid-base indicator is a weak acid (or a weak base) that has
different colors in its un-ionized and ionized states. Most indicators
are aromatic molecules that have delocalized electrons, and in Chapter
9 we saw the reason for their color changes. The equilibrium
with HIn representing the acid form of the indicator compound,
is shifted to the left by an excess of acid, and to the right by
an excess of base. The ratio of basic to acidic form of the indicator
is linked to the pH by the now familiar expression
in which pKa is the acid-dissociation constant for the
weak indicator acid, HIn. The eye is sensitive to color changes
over approximately a 1:10 to 10:1 concentration ratio, meaning that
visible color changes in an indicator occur in a pH range of around
2 units, centered on the indicator's own pKa. Litmus paper changes
from red in acid to blue in base, in the pH range 5-8. Phenolphthalein
solution added in minute quantities to the solution being tested
or titrated changes from colorless (acid) to red (base) in the range
of pH 8-10 because it has a pKa around 9. Other common indicators,
color changes, and useful pH ranges are shown in the diagram on
page 30.
SOLUBILITY EQUILIBRIA
The most familiar salt, sodium chloride, is so soluble that we
tend to think that all salts are equally soluble. This is far from
true, and many salts are quite insoluble in water.
Solubility is the result of competition between the mutual attractions
of ions in the crystal, and hydration of individual ions by solvent
molecules. Both of these processes involve large energies, and solubility
depends on the frequently quite small difference between them. It
is difficult to calculate crystal-lattice energies and hydration
energies accurately enough to predict whether the difference between
them will be positive or negative. Although we understand the forces
at work when a salt dissolves, it is not easy to predict whether
a given salt will dissolve or not.
There are a few common-sense principles that help in a general
way. A crystal is held together by electrostatic forces between
oppositely charged ions. Crystals with small ions that can be packed
close together generally are harder to pull apart than crystals
with large ions.
Hence fluorides (F-) and hydroxides (OH-)
tend to be less soluble than nitrates (NO3-)
and perchlorates (ClO4-) with the same positive
ion; chlorides (Cl-) are intermediate in solubility.