Titrations, Indicators and Titration Curves

The technique of titration is used to find out accurately how much of a chemical substance is dissolved in a given volume of a solution, that is, the concentration of the solution.

The technique uses a set of apparatus with which volumes of solutions can be measured to an accuracy of greater than 0.1 cm³. Three important pieces of apparatus are:

In a titration the pipette is used to transfer 25 cm³ (usually to ±0.05 cm³) of a solution into a conical flask. Another solution that reacts with the solution in the conical flask is carefully added from a burette until it has all exactly reacted. This is called the end point of the titration (or equivalence point of the reaction). There needs to be a way of knowing when the end point is reached. An indicator may be needed. Often a titration is repeated until successive titres are within 0.1 cm³.

An indicator is a substance that undergoes a change in colour when the end-point of a titration is reached. Acid-base indicators are used to signal the end of acid-base titrations.

Acid-base indicators are perhaps the most common types, but different types of indicators are used in precipitation reactions, such as in silver nitrate(V) titrations for chloride ion determination. In reactions where there is a colour change an indicator may not be needed, as in manganate(VII) titrations.

An acid-base indicator is itself a weak acid (or its conjugate base).

An acid-base indicator is a weak acid having a different colour in aqueous solution from its conjugate base.

Consider methyl orange, if the acid form of the indicator is represented by HIn and its conjugate base form by In^-, the following equilibrium exists in aqueous solution:

According to LeChatelier's Principle, the addition of an acid shifts the equilibrium to the left and the solution turns red. The addition of base removes H⁺, shifting the equilibrium to the right and the solution turns yellow.

The equilibrium condition for the reaction is:

Rearranging this expression:

Therefore, the ratio [HIn] / [In^-] depends on the pH, and determines the colour of the solution. With methyl orange, the solution is red if [HIn]>> [In^-], yellow if [In^-]>>[HIn], and varying shades of orange when [HIn] and [In^-] are about the same.

Therefore, at the end-point of the titration [HIn(aq)] / [In^-(aq)] » 1, and

K_a = [H₃O⁺(aq)]_eqm    or    pK_a = pH

pK_a for an indicator is (about) equal to the pH of the solution at the end-point.

Methyl orange as an indicator in strong acid-weak base titrations. It changes from red (at pH 3.1) to orange-yellow (at pH 4.4).

We can see the colour changes of methyl orange because it absorbs light in the visible part of the electromagnetic spectrum. Its molecule contains an extended system of delocalised electrons called a chromophore. The differences in energy between the quantised electronic energy levels correspond to the energies of photons of visible light. Electrons are promoted when these photons are absorbed, removing their frequencies from those that enter the eye. In methyl orange, when the molecule becomes protonated in acidic solution, the differences in energy between the electron energy levels change slightly from the unprotonated form. This results in the absorption of different frequencies of visble light and so a change in colour of the indicator. Methyl orange in acidic solution absorbs blue-green light, which makes its solution appear red. In alkaline solution it absorbs blue-green and red light making it appear yellow.

Titration curves

Acid-base indicators take advantage of the rapid change in pH of the solution being titrated as the equivalence point is reached. When an acid and base have been mixed in equivalent amounts (according to the chemical equation for the reaction) they are said to have neutralised each other. However, this term is somewhat misleading because the pH of the solution depends on the salt formed, and may not be pH 7.

The choice of an indicator is determined by the pH of the solution at the equivalence point.

For example, at the equivalence point of a titration involving ethanoic acid and sodium hydroxide, the only product is an aqueous solution of the ionic compound sodium ethanoate. It is the ethanoate ions behaving as a base that cause the solution at the end-point to have an alkaline pH.

CH₃COOH(aq) + OH^-(aq)   ®   H₂O(l) + CH₃COO^-(aq)

A studying of Ionic Equilibrium involving weak acids and their conjugate bases is needed to appreciate these ideas more fully.

Indeed, the pH of a solution formed at the equivalence point is important because it influences the choice of acid-base indicator for the titration. This is because...

acid-base indicators change colour within characteristic pH ranges.

Two familiar acid-base indicators are methyl orange and phenolphthalein.

Explore these ideas below:

25 cm³ of 0.1 mol dm^-3 acid is titrated with 0.1 mol dm^-3 alkaline solution.

Strong Acid - Strong Base
Hydrochloric acid, HCl(aq) - Sodium hydroxide, NaOH(aq)

Strong Acid - Weak Base
Hydrochloric acid, HCl(aq) - Ammonia, NH₃(aq)

Weak Acid - Strong Base
Ethanoic acid, CH₃COOH(aq) - Sodium hydroxide, NaOH(aq)

Weak Acid - Weak Base
Ethanoic acid, CH₃COOH(aq) - Ammonia, NH₃(aq)

Phenolphthalein
pH 8.3 - 10.0 Methyl Orange
pH 3.1 - 4.4 Phenolphthalein
pH 8.3 - 10.0 Methyl Orange
pH 3.1 - 4.4

Any acid-base indicator that changes colour between pH 4 and pH 10 is suitable to detect the end-point for a strong acid - strong base titration. Both methyl orange and phenolphthalein could be used. Just one drop of the added base will bring about a change in colour of the indicator.

The pH curve for the strong acid - weak base titration shows that phenolphthalein is not a suitable indicator but methyl orange is fine. For the titration of weak acid - strong base, phenolphthalein, but not methyl orange, is a suitable indicator. For a weak acid - weak base titration, the pH curve shows there is no rapid change in pH corresponding to the addition of just one or two drops of the base. For this reason it is not usually possible to detect the end-point using an acid-base indicator.