This indicator cannot, however, be successfully used with weak bases,
even ammonium hydroxide; for, since it is weak acid, the salts
which it forms with weak alkalies are easily hydrolyzed, and as a
consequence of this hydrolysis the change of color is not sharp.
This indicator can, however, be successfully used with strong bases,
because the salts which it forms with such bases are much less
hydrolyzed and because the excess of OH^{-} ions from these bases also
diminishes the hydrolytic action of water.

This indicator is affected by even so weak an acid as carbonic acid,
which must be removed by boiling the solution before titration. It is
the indicator most generally employed for the titration of organic
acids.

In general, it may be stated that when a strong acid, such as
hydrochloric, sulphuric or nitric acid, is titrated against a strong
base, such as sodium hydroxide, potassium hydroxide, or barium
hydroxide, any of these indicators may be used, since very little
hydrolysis ensues. It has been noted above that the color change does
not occur exactly at theoretical neutrality, from which it follows
that no two indicators will show exactly the same end-point when acids
and alkalis are brought together. It is plain, therefore, that the
same indicator must be employed for both standardization and analysis,
and that, if this is done, accurate results are obtainable.

The following table (Note 1) illustrates the variations in the volume
of an alkali solution (tenth-normal sodium hydroxide) required to
produce an alkaline end-point when run into 10 cc. of tenth-normal
sulphuric acid, diluted with 50 cc. of water, using five drops of each
of the different indicator solutions.