What is an Indicator?

Indicators are weak acids or bases, with a conjugate acid/base of a different colour. They will change colour in response to changes in the pH of a solution. As a weak conjugate acid-base pair, indicators form an equilibrium system, forming the equation:

When an acid (low pH) is added to the system, this increases the hydronium ion concentration and shifts the equilibrium to the left. This changes the colour of the solution towards Colour X. When a base (high pH) is added to the system, this decreases the number of hydronium ions, shifting the equilibrium towards the right. This changes the colour towards Colour Y

Equivalence Point and End Point

The ultimate goal of a titration is to end the titration at the ‘equivalence point’. This is the point in a titration where the exact amount of acid and base is present in the conical flask to neutralise each other. The colour changing properties of indicators can help us to approximate when the equivalence point has been reached. The point when the indicator changes colour is the ‘end point’. This end point occurs over a pH range, which varies for each indicator:

Titration Curve

A titration curve can help us to visualise the progress of a titration. A titration curve plots the pH of the acid or base being analysed (y-axis), against the volume of the acid or base with a known concentration being added (x-axis).

Each titration curve will follow a similar pattern. For example, consider a base (in the burette), is titrated against an acid (in the conical flask):

1. Initial slow pH change: gradual neutralisation of hydroxide and hydronium ions.

2. Steep pH change: all of the hydronium ions have been consumed by the hydroxide ions, i.e. the acid and base have been neutralised. This is the location of the equivalence point.

3. Final slow pH change: if even more base is added, there will be more hydronium ions than neutralised ions, shifting the pH towards the pH of the base. This is also known as overshooting the end point.

The end point of the chosen indicator should fall somewhere along the steep area of the curve where the equivalence point can be found.

Choosing an Appropriate Indicator

The closer the end point is to the equivalence point, the more accurate the titration will be. Therefore, the chosen indicator must have a pH range which roughly corresponds to the pH range of the neutralised solution. This can be determined by the salt formed by the acid and base used in the titration.

Depending on the strengths of the acid and base used in the titration, the equivalence point can be slightly acidic, basic, or neutral. This is because of the nature of the salt formed by the neutralisation of the acid and base.

Strong acids and bases will produce neutral ions. Weak acids will produce basic ions. Weak bases will produce acidic ions.

When a strong acid and base is used, they will produce a neutral ion. Strong acids and bases ionise completely. Therefore, their conjugate ions will not accept or donate protons to form back into their acid or base (e.g. sodium chloride).

Weak acids and bases only partially ionise. When a weak acid is used it will produce a conjugate ion that is basic (e.g. acetate). This will cause the equivalence point to be basic.

When a weak base is used it will produce a conjugate ion that is acidic (e.g. ammonium chloride). This will cause the equivalence point to be acidic.

Weak acids and bases are generally not used together during a titration. On a titration curve, there is no clear steep change in pH close to the equivalence point. Therefore, an appropriate indicator cannot be determined.