Carry out two redox reactions and observe and interpret the results in this class practical

In this experiment, students observe and interpret two redox reactions using microscale apparatus. First, students investigate the reaction between copper(II) ions and halide ions, adding drops of copper(II) sulfate solution to sodium chloride, potassium bromide and potassium iodide. They then explore the reaction between silver(I) ions and iron(II) irons using silver nitrate solution and iron(II) solution.

The practical should take approximately 20 minutes.



  • Student worksheet
  • Clear plastic sheet (eg ohp sheet)
  • Magnifying glass



Solutions should be contained in plastic pipettes – see the accompanying guidance on apparatus and techniques for microscale chemistry.

  • Potassium bromide, 0.2 mol dm–3
  • Potassium iodide, 0.2 mol dm–3
  • Sodium chloride, 0.2 mol dm–3
  • Silver nitrate, 0.1 mol dm–3
  • Copper(II) sulfate, 0.2 mol dm–3
  • Iron(II) sulfate, 0.2 mol dm–3
  • Iron(III) nitrate, 0.2 mol dm–3
  • Potassium thiocyanate, 0.1 mol dm–3
  • Starch solution (freshly made)

Health, safety and technical notes

  • Read our standard health and safety guidance.
  • Wear eye protection throughout (splash-resistant goggles to BS EN166 3).
  • Potassium bromide, KBr (aq), 0.2 mol dm–3 is low hazard.
  • Iron(II) sulphate, FeSO4.7H2O (aq), 0.2 mol dm–3 is low hazard.
  • Iron(III) nitrate, Fe(NO3)3.9H2O (aq), 0.2 mol dm–3 is low hazard.
  • Potassium thiocyanate, KBr (aq), 0.1 mol dm–3 is low hazard.
  • Potassium iodide, KI(aq), 0.2 mol dm–3 is low hazard.
  • Silver nitrate, AgNO3(aq), 0.1 mol dm–3 is an eye IRRITANT. Keep separate from organic waste containers.
  • Copper(II) sulphate solution, CuSO4(aq), 0.2 mol dm–3 causes eye damage and is HAZARDOUS to the aquatic environment.


Part 1: the reaction between copper(II) ions and halide ions

  1. Cover table 1 on your worksheet with a clear plastic sheet.
  2. Put one drop of copper(II) sulphate solution in each of the boxes below.
  3. Add one drop of sodium chloride solution to the first box; one drop of potassium bromide solution to the second box; one drop of potassium iodide solution to the third box. Observe.
  4. Add one drop of starch solution to each of the reaction mixtures. Observe.

Part 2: the reaction between silver(I) ions and iron(II) ions

  1. Cover table 2 on your worksheet with a clear plastic sheet.
  2. Put one drop of silver nitrate solution in the box below.
  3. Add one drop of iron(II) solution. Observe closely. What happens?
  4. After one minute add one drop of thiocyanate solution.
  5. To help you interpret your observations, put one drop of potassium thiocyanate solution in each of the boxes in table 3 on your worksheet. Add one drop of each of the reagents indicated and observe.


What explanations can you give for your observations in each experiment?

Teaching notes and expected observations

Part 1

No changes are observed on adding chloride or bromide to the copper(II) solution. However, the addition of iodide gives an immediate light brown precipitate of copper(I) iodide. The addition of starch solution gives the intense blue-black colour characteristic of the starch–iodine complex. Iodide reduces copper(II):

2Cu2+(aq) + 4I(aq) → 2CuI(s) + I2(s)

Part 2

The addition of iron(II) solution to silver nitrate produces silver metal by reduction. Glittering can be seen in the drop.

The addition of a drop of thiocyanate produces a deep red colour indicative of iron(III). A whitish precipitate of silver thiocyanate can also be seen.

The second part of this experiment is for students to do sequential reactions of thiocyanate with silver(I), iron(II) and iron(III), helping them to interpret this redox reaction.


Unless very pure and freshly prepared, iron(II) solutions will contain a small amount of iron(III) which gives a slight red coloration in the reaction between the iron(III) solution and the thiocyanate. However, the intensity of the colour is less than that observed in the reaction between iron(III) solution and thiocyanate ions. This point could be explored further in subsequent discussions on the purity of chemicals.