Indicators – choosing and using

Indicators – choosing and using

An indicator is a chemical that indicates the end-point of a reaction by changing colour.

Types of indicators

Acid – base indicators

The most common class of these are acid/base indicators such as litmus or phenolphthalein. These change colour at a particular pH.

To be exact, they change colour over a range of pH values

They are generally weak acids

HIn ⇋ H⁺  In^–

Depending on the pH of the solution, the equilibrium will shift from one side to the other. This is true for all weak acids. For indicators, however, HIn, is a different colour than In^–.

Universal (Full range) indicators

These are special examples of acid-base indicators. They are mixtures of several single acid-base indicators which all change at different pH values.

As a result at any pH there will be different amounts of different colour species present, leading to a continuous range of colours from one end of the pH scale to the other.

Full range indicator is just a Universal indicator that covers the whole range from 1 – 14.

Redox indicators

Indicators can be used for purposes other than simply telling pH. In this case, the indicator undergoes a definite color change at a specific electrode potential rather than pH value.

Common examples include phenanthroline and dephenylamine as indicators in the determination of iron.

Complexometric indicators

These are compounds that undergo a definite color change in presence of specific metal ions.

They form a weak complex with the ions present in the solution, which have a significantly different color from the un-complexed form.

Some common examples are  eriochrome black T and murexide for EDTA titrations of calcium and magnesium.

Precipitation indicators

In this case, the indicator is a compound that will, at the end point, produce a recognisable precipitate, often of a different colour.

The most well known is the use of potassium chromate for the titration of chloride ions with silver nitrate. As soon as all the chloride ions have been used up, a red precipitate of silver chromate appears.

Self indicators

The most common example here is potassium manganate VII. The permanganate ion is coloured an intense purple so as it reacts the colour fades, thus it acts as its own indicator.

Choosing indicators

For acid-base indicators, you need to consider the neutralisation reaction you are carrying out and pick and indicator that will change in the over the mid-point of the reaction.

For redox indicators, you will need to consider the electrode potentials involved – though to be more realistic, you will be more likely to use an indicator suggested by the protocol

As far as complexometric indicators go, you need to determine which indicator will complex with the ions you are investigating.

Using indicators

Preparing indicators

Indicators are usually intensely coloured substances so solutions are generally dilute – around 1% is a typical working concentration.

When added to the solution, it is usual to use no more than a few drops – enough to give a definite colour change.

Using too much indicator can interfere with the reaction. Adding a weak acid (an indicator) to an acid-base titration could make a difference if the solutions were very weak.

Some indicators are added in solid form – for instance it is common to grind murexide with potassium chloride (still in a 1% composition) and add a small amount of the solid to the reaction mixture.

Storing indicators

Most solid indicators will last for years.

Many solutions have an extremely good shelf life as well – especially the acid-base ones. Some, however, are more fragile and need to be made up fresh each time.