Electrolytic cells

Introduction

Electrolysis lends itself to some good practical work. Unfortunately there is little on the syllabus for Standard Level students to do practically. Sub-topic  9.2 Electrochemical cells requires students to be able to state what happens at each electrode and to describe how current is conducted but the examples are limited to the electrolysis of a molten salt. The Additional Higher Level Material takes this further in 19.1 Electrochemical cells (AHL) where electrolysis in aqueous solution is included as is the use of electrolysis for electroplating. Even for the AHL quantitative problems involving Faraday’s Laws do not appear to be required as only relative amounts of products formed during electrolysis. need to be determined. However Faraday's constant is still needed to solve problems using ΔG⦵ = − nFE⦵.

This practical includes the electrolysis of molten lead(II) bromide which covers the core but much of the practicals is concerned with covering the AHL as it gives examples in aqueous solution. I think it is reasonable to also expose SL students to this as it reinforces the idea of how current is conducted. However for SL students I would not stress too much the factors affecting which ions will gain or lose electrons in aqueous solution or the factors affecting which ions will be discharged.

Teacher's notes
 

There are several separate experiments. As they are not assessed I get the students to work in pairs in a ‘circus’ so they do one experiment then move on swapping experiments with other pairs so that eventually all the pairs of students have done each experiment.

Lead(II) bromide is chosen as the salt to melt as its melting point is much lower than the melting point of sodium chloride. (The melting point of PbBr2 is usually taken as 370 oC compared to the melting point of NaCl which is 801 oC). I usually get the students to put some of the white solid lead(II) bromide into a crucible and connect the circuit up before heating. All this should be done in a fume cupboard. They can check that the circuit is working by shorting across the electrodes. As the solid begins to melt they should see the reading on the ammeter start to increase and see bubbles of a gas coming from the positive electrode. Don’t let them heat it for too long after it has started conducting to minimise the fumes of bromine. They should observe the current begin to decrease and eventually return to zero as the salt resolidifies.

The electrolysis of the various solutions should be reasonably straightforward. The electrolysis of copper(II) sulfate using copper and/or graphite electrodes nicely illustrates the effect of changing the nature of the electrodes and the electrolysis of sodium chloride solution illustrates the effect of concentration. If you have a Hofman voltameter (see photograph on right) then it is well worth using it for the electrolysis of ‘water’ (dilute sulfuric acid or dilute sodium hydroxide solution) as it shows that exactly twice the volume of hydrogen gas is evolved compared to the volume of oxygen gas evolved. (This needs to run for a little while before it is exactly true as initially some of the oxygen evolved will dissolve in the dilute sulfuric acid until it becomes saturated). You can easily also show this using much simpler apparatus and just collecting the gases by displacing water from inverted test-tubes. 

Standard Level Higher Level Student worksheet

ELECTROLYTIC CELLS

Solid salts do not conduct electricity as their ions are in fixed positions and they contain no free electrons. When in the molten state or in aqueous solution they do conduct electricity as their ions are free to move. These ions are oxidised or reduced at their respective electrodes. The products that are evolved at the electrodes can depend upon a number of factors. The following qualitative experiments illustrate some of these factors.

ENVIRONMENTAL CARE:

This practical involves five different electrolytic reactions. To minimise the use of chemicals move round each experiment in turn and use the same solutions or solid as the previous users. At the end of the practical the copper(II) sulfate solution can be filtered and returned to the marked bottle, and any lead salt remaining should be placed in the 'Heavy Metal Waste' container. The remaining solutions can be disposed of down the sink.

SAFETY:

Take care when heating the lead(II) bromide. In some of the experiments small quantities of poisonous gases will be produced. Once you have identified these, switch off the current.

PROCEDURE:

Using a suitable circuit with a potential difference of between 2 and 4 volts pass a current of between 0.1 and 1.0 amps through the following electrolytes. Observe carefully all the changes taking place at both the positive electrode and the negative electrode and also observe whether there is any obvious change to the electrolyte itself. If gases are evolved try to identify them and note whether they are evolved at the positive or negative electrode. If gases are evolved at both electrodes observe whether they are given off in similar volumes.

1. Copper(II) sulfate solution. Using two copper electrodes allow the current to flow for five minutes. Repeat the experiment using two graphite electrodes.

2. Molten lead(II) bromide. Connect the circuit whilst the lead(II) bromide is still solid. Heat gently in order to melt the solid and allow the current to flow for a few minutes. Take care not to let the Bunsen flame melt the connecting wires!

3. Sodium chloride solution. Using graphite electrodes pass the current through a dilute solution of sodium chloride for a few minutes then repeat the experiment using the more concentrated solution. After each experiment add a few drops of indicator solution to the electrolyte near each electrode.

4. Potassium iodide solution. Use graphite electrodes and allow the current to flow for a few minutes.

5. Water. Try to electrolyse distilled water using graphite electrodes. If nothing happens add a few drops of dilute sulfuric acid to the water.

QUESTIONS

1. Try to explain all your observations giving the relevant half-equations.

2. List the factors that can affect the discharge of ions during electrolysis and give one example of each from the above reactions.

3. Describe how electrical conductivity by molten or aqueous salts is different to electrical conductivity by a metal.

This worksheet can also be downloaded from:

  Electrolytic cells  

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