Acid-Base Indicators

Many substances, including that one almost every one associates with acids and bases – litmus – change color in response to acid or base. The pigment in red cabbage is a natural substance very commonly used to show color change. Phenolphthalein is one of the most common indicators used for beginning chemistry, because its color change is very obvious which makes it easy to use.

Different dyes will change color at different pH's (the value can be calculated from the equilibrium constant for the indicator). Here is a small sample of some common acid-base indicators, and the range at which their pH changes color.

One of the difficulties with giving a range of colors is that different person's eyes are not all equally sensitive (also different monitors will display colors differently), so these colors are only approximations.

When doing a titration, the color of the dye will change at the endpoint. The equivalence point is the point at which the amount of H₃O⁺ and OH^- are equal. When conducting a titration, one must select the proper indicator so that the pH at its endpoint will match the equivalence point of the titration. There is a rapid change in pH as you approach the equivalence point. By measuring or calculating the titration curve, you can determine which indicator will work best.

Acid-Base indicators are dyes that are themselves weak acids and bases. However, the conjugate acid-base forms of the dye have different colors. The actual chemical structures of the dyes is often quite complex; however, we use can the generic symbol for the indicator as HIn. The Brönsted-Lowry equation for the indicator is:

HIn + H₂O H₃O⁺ + In^-

where the color of the letters is used to show the differently colored forms of the dyes, and the sizes of the letters the relative concentrations.

Suppose that we increase the concentration of H₃O⁺ in a more acidic solution. Then if we apply le Châtelier's principle, the result should be

HIn + H₂O H₃O⁺ + In^-

There will be more H₃O⁺ (because we added more), but the system will respond to:

use up some of the added H₃O⁺, at the same time
removing some of the In^-, and
making more HIn.

Thus we will see the color of the HIn form.

Now suppose that we decrease the concentration of H₃O⁺(which we could do by adding more OH^- in a more basic solution). Applying LeChâtelier's principle, the result should be

HIn + H₂O H₃O⁺ + In^-

There will be less H₃O⁺ (because we removed it), but the system will respond to try and:

make some more H₃O⁺, at the same time
making more In^-, and
removing some of the HIn.

Thus we will see the color of the In^- form.

In summary: at a low pH, an indicator is almost entirely in the HIn form. As the pH increases, the intensity of the colour of In^- increases as the equilibrium shifts to the right.