A. Acid-base indicators

Acid-base indicators are weak organic acids or weak organic bases that indicate whether a solution is acidic, basic or neutral.

Let us consider a weak organic acid that is denoted by HIn. In order to be effective indicators, HIn, and its conjugate base, In–, must have different color. In solution, the acid ionizes as follows:

If the indicator is in sufficiently acidic medium, the equilibrium, according to Le Chatelier’s principle, shifts to the left and the predominant color of the indicator is that of non-ionized form (HIn). On the other hand, in a basic medium, the equilibrium shifts to the right and the color of the solution will be that of the ionized form (In^–).

This means that the concentration on HIn is 100 times that of In–, and so the solution appears yellow.

At pH of 6.0, [H₃O⁺] is 1.0 × 10^–6 equal concentrations of HIn and In^– give the solution an orange color.

At pH = 7.0, [H₃O⁺] is 1× 10^–7. Here the concentration of In^– is 10 times that of HIn, and so the solution appears red.

B. Acid-base titration

To determine the concentration of a particular solute in a solution, chemists often carry out a titration. Titrations can be conducted using neutralization, precipitation, or oxidation-reduction reactions.

A titration is a technique in which a solution of known concentration is used to determine the concentration of an unknown solution. Typically, the titrant (the known solution) is added from a burette to a known quantity of the analyte (the unknown solution) until the neutralization reaction is complete. The point at which the acid has completely reacted with or been neutralized by the base, or vice versa, is called the equivalence point of the titration.

According to the definition of normality, the number of equivalents is the normality multiplied by the volume of solution, in litres. If we add enough acid to neutralize a given volume of base, the following equation holds:

N_aV_a = N_bV_b

Where N_a and V_a refer to the normality, and volume of the acid solution, respectively, and N_b and V_b refer to the normality and volume of the base solution, respectively.

To perform a successful titration, we must use an indicator that changes color at the equivalence point.

Acid-base titration curves

An acid-base titration curve is a plot of the pH of a solution of acid (or base) against the volume of added base (or acid). Such curves are used to gain insight into the titration process. You can use the titration curve to choose an indicator that will show when the titration is complete.

Titration of a strong acid and a strong base

Features of the titration curve for the titration of a strong acid with a strong base.

• The pH is low at the beginning of the titration.
• The pH changes slowly until just before the equivalence point.
• Just before the equivalence point, the pH rises sharply.
• At the equivalence point, the pH is 7.0.
• Just past the equivalence point, the pH continues to rise sharply.
• Further beyond the equivalence point, the pH continues to increase, but much less slowly.

Titration of a weak acid and a strong base

In contrast to the titration of a strong acid with a strong base, the titration of a weak acid with a strong base, has these features:

• The initial pH is higher because the weak acid is only partially ionized.
• At the half-neutralization, pH = pK_a. The solution at this point is a buffer solution in which the concentration of the weak acid and its conjugate base are equal.
• The pH is greater than 7 at the equivalence point because the anion of the weak acid hydrolyzes.

Titrations of a weak base and a strong acid

How does the titration curve of a weak base with a strong acid differ from the titration curves you have seen so far? What is the pH at equivalence point?

Consider the titration of ammonia, NH3, with a strong acid, HCl.

HCl (aq) + NH₃ (aq) ⟶ NH₄Cl (aq)

or simply

H⁺ (aq) + NH₃ (aq) ⟶ NH₄⁺ (aq)

The pH at the equivalence point is less than 7. Why?