should, then, be formed in the undissociated condition as soon as even
a slight excess of OH^{-} ions is present in the solution, and there
should be a prompt change from red to yellow as outlined above. On the
other hand, it should be an unsatisfactory indicator for use with weak
acids (acetic acid, for example) because the salts which it forms
with such acids are, like all salts of that type, hydrolyzed to a
considerable extent. This hydrolytic change is illustrated by the
equation:

(M.o.)^{+} C_{2}H_{3}O_{2}^{-} + H^{+}, OH^{-} --> [M.o.OH] + H^{+},
C_{2}H_{3}O_{2}^{-}.

Comparison of this equation with that on page 30 will make it plain
that hydrolysis is just the reverse of neutralization and must,
accordingly, interfere with it. Salts of methyl orange with weak acids
are so far hydrolyzed that the end-point is uncertain, and methyl
orange cannot be used in the titration of such acids, while with
the very weak acids, such as carbonic acid or hydrogen sulphide
(hydrosulphuric acid), the salts formed with methyl orange are, in
effect, completely hydrolyzed (i.e., no neutralization occurs), and
methyl orange is accordingly scarcely affected by these acids. This
explains its usefulness, as referred to later, for the titration of
strong acids, such as hydrochloric acid, even in the presence of
carbonates or sulphides in solution.

Phenolphthalein, on the other hand, should be, as it is, the best of
the common indicators for use with weak acids. For, since it is
itself a weak acid, it is very little dissociated, and its nearly
undissociated, colorless molecules are promptly formed as soon as
there is any free acid (that is, free H^{+} ions) in the solution.