Theory
In and of themselves, pH indicators are frequently weak acids or weak bases. The general reaction scheme of a pH indicator can be formulated as follows:
- HInd + H2O H3O+ + Ind-
Here HInd stands for the acid form and Ind- for the conjugate base of the indicator. It is the ratio of these that determines the color of the solution and that connects the color to the pH value. For pH indicators that are weak protolytes, we can write the Henderson-Hasselbalch equation for them:
The equation, derived from the acidity constant, states that when pH equals the pKa value of the indicator, both species are present in 1:1 ratio. If pH is above the pKa value, the concentration of the conjugate base is greater than the concentration of the acid, and the color associated with the conjugate base dominates. If pH is below the pKa value, the converse is true.
Usually, the color change is not instantaneous at the pKa value, but there is a pH range where a mixture of colors is present. This pH range varies between indicators, but as a rule of thumb, it falls between the pKa value plus or minus one. This assumes that solutions retain their color as long as at least 10% of the other species persists. For example, if the concentration of the conjugate base is ten times greater than the concentration of the acid, their ratio is 10:1, and consequently the pH is pKa + 1. Conversely, if there is a tenfold excess of the acid with respect to the base, the ratio is 1:10 and the pH is pKa – 1.
For optimal accuracy, the color difference between the two species should be as clear as possible, and the narrower the pH range of the color change the better. In some indicators, such as phenolphthalein, one of the species is colorless, whereas in other indicators, such as methyl red, both species confer a color. While pH indicators work efficiently at their designated pH range, they are usually destroyed at the extreme ends of the pH scale due to undesired side-reactions.
Read more about this topic: PH Indicator
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