Transition metals: Using a colorimeter to measure coloured ion concentration

Edexcel A level Chemistry (2017)

Topic 15A: Principles of transition metal chemistry

Using the Beer-Lambert Law to measure the unknown concentration of a transition metal ion in solution.

Colorimetry is used to measure the concentration of a coloured ion in solution.

In Colorimetry, we measure how much visible light a coloured solution of a given concentration is absorbing. 

The measurements are taken using a colorimeter. (see diagram)

The light passes through a filter.  The filter selects the wavelength of light closest to maximum absorption. The coloured solution (in a cuvette) absorbs some of this light. 

The light absorbed is called the absorbance.  The greater the concentration of coloured ions, the greater the absorbance.

In less expensive colorimeters like those you have in school, you first zero the instrument using a cuvette of 'colorless' water. 

This sets readings to zero absorbance.

Next you put in a cuvette of the coloured solution and read off the 'absorbance'.

‘Zeroing' eliminates error because even a 'colorless blank' of cuvette and water can absorb a tiny amount of light.

How you can use colorimetry to measure the concentration of a transition metal ion in solution.

You will need to create a calibration curve of absorbance vs standard concentrations and use it to measure the concentration of an unknown solution of your transition metal ion. 

The absorbance of an intensely coloured transition metal ion in aqueous solution can be measured directly e.g. the concentration of manganese in the deep purple manganate(VII) ion, MnO₄^—.

But the absorbances of much less intensely coloured solutions of transition metals are more difficult to measure e.g. light blue hexaaquacopper(II) ions.

So the thing to do is to create an intensely coloured solution using ligand exchange. The intensely coloured solution then gives accurate measures of absorbance that can then be turned into accurate measures of concentration. 

So add ammonia solution to the blue hexaaquacopper(II) ion to form the deep blue ammine complex.

Or add potassium thiocyanate to a brown solution of iron(III) ions to form the deep blood–red thiocyanate ion (SCN^–) complex.

Creating the calibration curve

To do this set up a range of concentrations of your transition metal ion solution.  Next you have to measure the absorbance of each of these solutions, preferably more than once, to establish a reliable calibration curve. 

Next plot your results (see below) This plot shows results for the absorbance of different concentrations of copper (II) ions. 

Providing you work at reasonably low concentrations (probabaly <1.0M), you will find that the calibration curve is a straight line (as above).  In other words, it obeys a law, the Beer—Lambert Law: as we said above, the greater the concentration, the greater the absorbance. 

Absorbance is directly proportional to the concentration of the aqueous solution. 

All you have to do now is take your solution of transition metal ion whose concentration you don’t know, measure its absorbance and read off the concentration from the calibration curve.