The illustration below shows the relationship between pH and an indicator’s color. The ladder diagram defines pH values where each of the indicator’s two forms
- the weak acid form, HIn, which in this example is yellow
- the weak base form, In–, which in this example is red
are the predominate species. The indicator changes color when the pH is between pKa – 1 and pKa + 1, changing from yellow to orange to red (or from red to orange to yellow).
The photo here shows the actual colors for this indicator, which is methyl red, in its (a) weak acid, HIn, form, (c) its weak base, In–, form, and (b) a mixture of HIn and In–.
The relatively broad range of pH levels over which an indicator changes color means that using a visual indicator to locate the titration’s end point is subject to uncertainty. To minimize a determinate titration error, an indicator’s entire pH range must fall within the rapid change in pH that occurs near the titration’s equivalence point. As shown in the figure below, phenolphthalein is an appropriate indicator for the titration of 50.0 mL of 0.050 M acetic acid with 0.10 M NaOH. Bromothymol blue, on the other hand, is an inappropriate indicator because its change in color begins before the initial sharp rise in pH, and, as a result, spans a relatively large range of volumes. The early change in color increases the probability of obtaining inaccurate results, while the range of possible end point volumes increases the probability of obtaining imprecise results.
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