ACID–BASE INDICATORS
In Section 11-2 we described acid–base titrations and the use of indicators to tell us when
to stop a titration. Detection of the end point in an acid–base titration is only one of the
important uses of indicators.
An indicator is an organic dye; its color depends on the concentration of H 3 Oions,
or pH, in the solution. By the color an indicator displays, it “indicates” the acidity or
basicity of a solution. Figure 19-1 displays solutions that contain three common indica-
tors in solutions over the pH range 3 to 11. Carefully study Figure 19-1 and its legend.
The first indicators used were vegetable dyes. Litmus is a familiar example. Most of
the indicators that we use in the laboratory today are synthetic compounds; that is, they
have been made in laboratories by chemists. Phenolphthalein is the most common acid–
base indicator. It is colorless in solutions of pH less than 8 ([H 3 O] 10 ^8 M) and turns
bright pink as pH approaches 10.
Many acid–base indicators are weak organic acids, HIn, where “In” represents various
complex organic groups. Bromthymol blue is such an indicator. Its ionization constant is19-4
808 CHAPTER 19: Ionic Equilibria II: Buffers and Titration Curves
Phenolphthalein was the active
component of the laxative Ex-Lax. It is
sometimes added to laboratory ethyl
alcohol to discourage consumption.
Figure 19-1 Three common
indicators in solutions that cover
the pH range 3 to 11 (the black
numbers). (a) Methyl red is red
at pH 4 and below; it is yellow at
pH 7 and above. Between pH 4 and
pH 7 it changes from red to red-
orange, to orange, to yellow.
(b) Bromthymol blue is yellow at
pH 6 and below; it is blue at pH 8
and above. Between pH 6 and 8 it
changes from yellow to yellow-
green, to green, to blue-green, to
blue. (c) Phenolphthalein is colorless
below pH 8 and bright pink above
pH 10. It changes from colorless to
pale pink, to pink, to bright pink in
the pH range 8 to 10.
34567 891011(a)345 67 8 91011(b)34567 891011(c)