You will learn more details about how indicators function to inform us that we have added sufficient acid or base to counteract another base or acid.
The goal of a titration is to add an exact stoichiometric equivalent of one substance to another with which it is reacting. One way to estimate the equivalence point in an acid-base titration is to use an indicator and observe the color change. The indicator is a dye of some sort and the color change occurs at the end point of the titration. We try to make the difference between the end point and the equivalence point negligible. The dye is usually an organic acid, Hind, with the acid having one color and the base, Ind-, having another color:
Hind(aq) + H2O(l) Ind-(aq) + H3O+(aq)
The equilibrium constant for the indicator will be Ka = [Ind-][H3O+]/[Hind]. You see that when the hydronium ion concentration in the solution is high, the [Hind] will be high and so the color will be that of the acid form of the indicator.
From the equilibrium constant equation we see that when [H3O+] is equal to the Ka, then [Ind-] will equal [Hind]. From this relationship, you might think that we could easily see the color change at this equality. However our vision is not quite that good. Most people can see the color of Hind when [Hind]/[Ind-] is about 10/1 and we can see the color of Ind- when [Hind]/[Ind-] is about 1/10. Thus most people observe the color change over a hydronium concentration interval of about 100, which corresponds to 2 pH units. From the last section we remember that the pH changes by as many as 7 units on passing through the equivalence point in a titration, so indicators can be very effectively used for titrations. From Figure 17.11, you see many indicators available over a wide range of effective pH units.
Returning to the equilibrium constant equation, Ka = [Ind-][H3O+]/[Hind], we see that we can rewrite it a
[H3O] = Ka[Hind]/[Ind-] which gives pH = pKa + log [Ind-]/[Hind] and if the color change occurs with a ratio of about 1/10 for [Ind-]/[Hind], then we have pH = pKa - 1, which is a useful equation to help find a suitable indicator for a particular titration.
Use Figure 17.11 to find an indicator to use in each of the following titrations
Web Author: Dr. Leon L. Combs
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