Electrochemical cells
9.2 Electrochemical cells (4 hours)
Pause for thought
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 students added metals to solutions containing the ions of other metals then they would have been able to deduce an activity series experimentally. What they probably did not notice was that in the cases where a reaction occurs the temperature of the solution increased slightly as the reactions are exothermic. They should however already have known this from when they carried out the mandatory practical for Topic 5 : Energetics/thermochemistry, such as to determine the enthalpy change for the reaction between zinc metal and copper(II sulfate solution. Since they now 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.
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 students often 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 to students since they have usually 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 students use the cell convention for voltaic cells. This requires them to 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.
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 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.
One small point is worth giving consideration to. Students 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 students needed to know which represented the positive pole of the battery (long thin line) and which the negative (short fat line). If they got this the wrong way round in the diagram they would 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
Nature of Science
The constant search to find different, more efficient and more environmentally friendly sources of energy is driven both by profit and by social needs. This is a good example of the ethical implications of science.
Learning outcomesAfter studying this topic students should be able to: Understand:
Voltaic cells:
Electrolytic cells:
Apply their knowledge to:
| Clarification notesA cell diagram convention should be used for voltaic cells. International-mindednessEnergy factors are a key component of space exploration. Hydrogen–oxygen fuel cells are used as an energy source in spacecraft. The International Space Station provides a good example of a multinational project involving the international scientific community. |
Teaching tipsFor voltaic cells I think is one of the few cases where I find that an initial demonstration by the teacher can be useful as you can take them through a simple cell stepwise. Use a big high resistance voltmeter and build up the system asking questions as you go so that you can point out all the aspects of a cell. For example asking them what the voltage reading will be before you add the salt bridge and then what they think will happen when different half-cells are connected so that you can build up your own 'electrochemical series' Use Mg(s)/Mg2+(aq), Zn(s)/Zn2+(aq), Fe(s)/Fe2+(aq) and Cu(s)/Cu2+(aq) as typical half-cells. You can also show them why it is important that when they draw a diagram they do not show the connection to the electrodes being in the solution as this introduces a new metal into the system. This should help them avoid this common mistake made by students when they are asked to draw a cell diagram in an exam. Then get them to do the practical Voltaic cells listed on the right which goes a little further and also covers the Mandatory laboratory component for this topic . It shows that the actual voltage depends not only on the nature of the two half-cells but also on the concentration of their ions. This can then help reinforce the concept of equilibrium as they can predict the effect on the voltage if the concentration of the ions is changed. Most students who have done some sort of pre-IB course will already have come across simple electrolysis. You should also have already mentioned the properties of ionic salts when you were covering Topic 4 : Chemical bonding and structure so it should not really be new to them. Reinforce that negative ions are called anions as they are attracted to the anode (positive electrode) in electrolysis where they become oxidized and positive ions are called cations as they are attracted to the cathode (negative electrode) in electrolysis where they become reduced. and that this is the opposite to voltaic cells The products from the electrolysis of simple molten salts are obvious as the metal is formed and either the halogen or oxide as the only salts that tend to be under consideration are metal halides and oxides. I have yet to see a question which asks what the products would be for the electrolysis of molten sodium sulfate or molten potassium carbonate. Finally stress how ionic salts are able to conduct in the molten state and emphasise that the conduction is due to the movement of ions through the molten liquid and that chemical decomposition must take place at the same time as electrons are gained at the cathode and lost at the anode. If your regulations allow the use of lead(II) bromide then it is good to set up the apparatus in a fume cupboard with an ammeter included in the circuit so that they can see that the solid does not conduct but as the salt starts to melt the current starts to flow. | Study guide
Page 73 QuestionsFor ten 'quiz' multiple choice questions with the answers explained see MC test: Electrochemical cells. For short-answer questions which can be set as an assignment for a test, homework or given for self study together with model answers see Electrochemical cells questions. Vocabulary list:Electrochemical cell IM, TOK, 'Utilization' etc.See separate page which covers all of Topic 9. Practical work |
Teaching slides
Teachers may wish to share these slides with students for learning or for reviewing key concepts.
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Other resources
1. A fairly straightforward but quite comprehensive description of how a simple voltaic cell works without going into standard electrode potentials.
2. A fairly thorough demonstration and description of the electrolysis of molten lead(II) bromide by Frank Scullion. It would of course be better to get your students to do this for themselves.
Electrolysis of molten lead(II) bromide
3. A nice simple practical by Bob Becker which is well-explained making a six-way galvanic cell.