Something to think about

This sub-topic follows on neatly from the last one on oxidation and reduction and in particular the activity series. If you did some practical work where you added metals to solutions containing the ions of other metals then you would have been able to deduce an activity series experimentally. What yu may not have noticed was that in the cases where a reaction occurs the temperature of the solution increased slightly as the reactions are exothermic. You should however already have known this from when you carried out the mandatory practical for  Topics 5 & 15 : Energetics/thermochemistry, such as to determine the enthalpy change for the reaction between zinc metal and copper(II sulfate solution. Since you understand that the reaction can be broken down into two half-equations involving electrons the challenge is to devise a method whereby the energy of the displacement redox reaction can be given out in the form of electrical energy instead of heat.

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The first voltaic cell was produced by Volta in 1792 but arguably the first reliable cell was the Daniel cell (above) which dates from 1836. This essentially consisted of a copper cylinder with the bottom sealed containing copper(II) sulfate solution. Inside this cylinder was another container. This was an earthenware pot (this allowed the transfer of ions) which was filled with sulfuric acid and a zinc electrode. The development of batteries has come a long way in the past two hundred years. However for the IB at Standard Level a simple 'voltaic cell’ can still be thought of as made up of two metals in solutions of their own ions connected by a salt-bridge and external wires.

There are several key points about voltaic cells that some students may find confusing.

• If a simple cell is connected to a high resistance voltmeter then the reading in volts is the cell’s potential energy or electromotive force (EMF). Until the cell does work (e.g. is used to power a light bulb) no current is flowing, no electrical energy is discharged and no reaction takes place.

• An electrochemical cell is the result of a spontaneous process. Unlike electrolysis (which is a non-spontaneous process) energy does not have to be put into the system.

• The negative electrode of the cell will be formed by the most reactive metal as it is giving up electrons. Since this is an oxidation process the negative electrode is sometimes called ‘the anode’. This can be extremely confusing since you may have come across ‘anode’ before in electrolysis where the anode is the positive electrode! For this reason many chemists prefer to use positive and negative electrodes rather than anode and cathode.
• The syllabus requires that you use the cell convention for voltaic cells. This means you shkould place the half-cell undergoing oxidation on the left hand side of the diagram and the half-cell undergoing reduction on the right hand side. The two aqueous solutions are then placed on either side of the salt bridge.

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Electrolysis is an important topic but there is very little about electrolytic cells on the syllabus at Standard Level as it only refers to molten electrolytes. Electrolytic cells are really the reverse of voltaic cells as non-spontaneous reactions are involved so energy in the form of electricity has to be supplied. No mention is made of electrolysis in aqueous solution so the factors affecting the discharge of ions are not relevant and practical applications such as electroplating and the refining of copper cannot be used. Arguably it does provide the theory for the production of aluminium which is part of Option A : Materials but even there the cryolite is added to make a solution not to lower the melting point so it is not really the electrolysis of molten aluminium oxide that is involved.

The other problem is that it is difficult for you to carry out practical work in a school laboratory on molten electrolytes. Simple salts such as sodium chloride or potassium bromide have melting points that are too high. The best examples to use are lead(II) bromide or lead(II) iodide but some schools are not happy with these as poisonous bromine or iodine gas is produced.

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One small point is worth giving consideration to. You are expected to be able to draw and annotate a diagram of an electrolytic cell. Traditionally a battery is represented by a 'long thin line and a short fat line'. This meant that you needed to know which represented the positive pole of the battery (long thin line) and which the negative (short fat line). If you get this the wrong way round in the diagram they will be penalised. Modern text books tend to physically draw a battery with labelled terminals so this problem does not arise. The IB does not seem to make it clear whether they would accept a battery with labelled terminals as part of the diagram in a written examination answer. Because they have not made it clear it would seem that either should be acceptable

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