The purpose of this experiment is to find out how a system in equilibrium responds to a change in concentration of components in the mixture.
Iron(III) ions and thiocyanate ions react in solution to produce thiocyanatoiron(III), a complex ion, according to the equation:
Fe3+(aq) + SCN–(aq) → Fe(SCN)2+(aq)
Pale yellow + colourless → blood-red
The colour produced by the complex ion can indicate the position of equilibrium.
- safety glasses
- 4 test-tubes and test-tube rack
- 2 teat-pipettes
- distilled water
- potassium thiocyanate solution, 0.5 M KSCN
- iron(III) chloride solution, 0.5 M FeCl3
- ammonium chloride, NH4Cl
- glass stirring rod
- Mix together one drop of 0.5 M iron(III) chloride solution and one drop of 0.5 M potassium thiocyanate solution in a test-tube and add about 5 cm3 of distilled water to form a pale orange-brown solution.
- Divide this solution into four equal parts in four test-tubes.
- Add one drop of 0.5 M iron(III) chloride to one test-tube. Add one drop of 0.5 M potassium thiocyanate to a second.
- Compare the colours of these solutions with the original samples. Record your observations.
- Add a spatula-full of solid ammonium chloride to a third test-tube and stir well. Compare the colour of this solution with the remaining tube and note your observation. Ammonium chloride removes iron(III) ions from the equilibrium by forming complex ions such as FeCl4–. A possible reaction is:
Fe3+(aq) + 4Cl–(aq) → FeCl4–(aq)
Interpretation of results
Having made three observations, suggest a cause for each colour change (in terms of the concentrations of the coloured species) and then suggest what can be inferred about a shift in the position of equilibrium. If a pattern has emerged, then you can make a prediction based on the results of the experiment.
- How would the position of equilibrium be affected by increasing the concentration of FeSCN2+?
- For each imposed change show how the shift in equilibrium position conforms to Le Chatelier’s principle.