Electrochemical cells (AHL)
19.1 Electrochemical cells (6 hours)
Pause for thought
1. Voltaic cells and standard electrode potentials
This topic causes students considerable difficulty when they come to answer questions on standard electrode potentials in Paper 2. It is also a topic that many teachers seem to find difficult - maybe there is some relationship between these two statements? The problem seems to stem from what is meant by "standard electrode potential". It could be that the IB data booklet does not help. For example, the lithium half-cell at the top of the electrochemical series in Section 24 of the data booklet appears as:
Because of the way it is written students (and possibly teachers too) confuse this with an energetics equation and think that if they write the equation the other way round then E will be equal to + 3.04 V. This is the wrong way of looking at standard electrode potentials. It is better to consider the half-cell as lithium metal in contact with lithium ions, i.e. Li(s)/Li+(aq), and when this is connected to a standard hydrogen electrode the minus sign means that electrons will flow from this half-cell to the hydrogen half-cell and the voltage will be 3.04 V. The electrochemical series is simply a list of what happens when all the half-cells in question are connected in turn to a hydrogen half-cell. The only concept that is key to remember is that when any two half-cells are connected together to give a spontaneous reaction the electrons will flow from the more negative half-cell to the more positive half-cell and the difference between the two values will be the EMF (cell potential) for that cell.
Let's look at one example to illustrate this. What will happen spontaneously when an Fe2+(aq)/Fe3+(aq) half-cell is connected to an iodine half-cell, ½I2(s)/I–(aq)?
To solve this problem draw the cell diagram and write the E values from the data booklet under the relevant half-cells:
½I2(s)/I−(aq) || Fe2+(aq)/Fe3+(aq)
+ 0.54 V + 0.77 V
The more negative (actually the least positive) is the iodine half-cell so the electrons will flow from this half-cell to the Fe2+(aq)/Fe3+(aq) half-cell. That is, the iodine half-cell will lose electrons and the iron ion half-cell will gain electrons so the two half-equations will be:
Fe3+(aq) + e– → Fe2+(aq)
2I–aq) → I2(s) + 2e–
So the overall spontaneous reaction is that iron(III) ions will oxidize iodide ions to iodine. The total cell potential (EMF) for this reaction will be 0.23 V.
2Fe3+(aq) + 2I–aq) → I2(s) + 2Fe2+(aq) Ecell = 0.23 V
A similar question could be asked a different way. Will iron(III) ions oxidise bromide ions to bromine? Solve it the same way. Write the two half-cells, look up their standard electrode potentials and then determine which will be the spontaneous reaction.
Fe2+(aq)/Fe3+(aq) || ½Br2(l)/Br–(aq)
+ 0.77 V + 1.07 V
Now the iron ion half-cell has the more negative (less positive) value so the electrons flow from this half-cell to the bromine half-cell. The half-equations for the spontaneous reactions become:
Fe2+(aq) → Fe3+(aq) + e–
Br2(l) + 2e– → 2Br–aq)
The overall equation for the spontaneous reaction will be:
2Fe2+(aq) + Br2(l) → 2Br–(aq) + 2Fe3+(aq) Ecell = 0.30 V
What this tells us is that iron(III) ions will not spontaneously oxidize bromide ions to bromine. In fact bromine will spontaneously oxidize iron(II) ions to iron(III) ions.
If all this is understood students will be able to answer any questions thrown at them on the use of standard electrode potentials. I have written a separate page under Areas of difficulty for Paper 2 which gives some more examples.
One question that students sometimes ask which you may find difficult to answer is how the standard electrode potential of very reactive metals like sodium can be determined. You might find it interesting to read a research paper written by Edgar Reynolds Smith and John Keenan Taylor in 1940 which details how they obtained the value for the sodium half-cell. Essentially they make use of an indirect method which depends on firstly measuring the potential of dilute sodium amalgam with respect to an aqueous solution of its ions and a reference electrode, and secondly, the difference in potential between the sodium amalgam and the pure metal when both are immersed in a non-aqueous conducting solution containing sodium ions. This non-aqueous solvent must not react with sodium and must dissolve the alkali salt sufficiently to form an electrolytic solution.
2. Electrolysis of aqueous solutions
This is quite a difficult topic to teach well. The trick is to understand that only certain examples are required to illustrate the various factors that affect the products (and their amounts) produced during electrolysis.
Position in the electrochemical series. This is almost completely limited to the position of metals and their ions and, of course, hydrogen and water. The obvious examples to include here are the electrolysis of dilute sodium hydroxide solution which behaves as 'water' (as hydrogen is evolved in preference to sodium) and the electrolysis of copper(II) sulfate solution with inert electrodes (where copper is deposited in preference to hydrogen ions gaining electrons). The position of negative ions is not really required by the IB as under normal conditions hydroxide ions will give up electrons in preference to other anions. The obvious example of this is the electrolysis of 'water' using dilute sulfuric acid. It is nice to show this using a Hofmann voltameter (right) although this particular piece of apparatus is not on the syllabus as such.
Nature of the electrode. For the most part electrodes made of graphite or platinum are inert and play no part in the products. The classic examples where they do affect the products are the use of copper electrodes in the electrolysis of copper(II) sulfate and the now almost obsolete mercury cell once used in industry where a flowing cathode of mercury results in a sodium amalgam product.
Concentration of the electrolyte. The classic example here is sodium chloride solution. For dilute solutions the products are almost exclusively hydrogen and oxygen but as the solution becomes more concentrated chlorine rather than oxygen is evolved at the positive electrode (anode) although a mixture containing both gases is actually formed. This is fertile ground for many a chemistry Extended Essay and could perhaps also be used for the individual scientific investigation.
Relative amounts of products. This is the application of Faraday's Laws although the laws themselves are not required. What this does mean is required is the concept of connecting two electrolysis cells in series. This way the same amount of electricity must be passing through both cells and if the amount of one product is known the amount of other products can be deduced from their relevant half-equations.
Electroplating. Examples here include the refining of copper using a small piece of pure copper as the negative electrode and impure copper the positive electrode (although the use of that lovely phrase 'anodic sludge' is not required) and plating with metals such as chromium where the object to be plated is made the negative electrode (cathode).
Nature of science
The use of standard electrode potentials, whose values are all relative to the standard hydrogen electrode, provides a good example of quantitative reasoning in science.
Scientists working together on electrochemical cell technologies, where they have had to consider the environmental and ethical implications of using fuel cells and microbial fuel cells, provides a good example of collaboration and ethical implications.
Learning outcomesAfter studying this topic students should be able to: Understand
Apply their knowledge to:
| Clarification notesThe theory of the electrolysis of aqueous solutions (e.g. sodium chloride, copper(II) sulfate and 'water' etc.) using both inert electrodes (platinum or graphite) and copper electrodes should be covered. The relevance of the E⦵ values, the nature of the electrode and the concentration of the electrolyte should be included in the explanations. The equation ΔG⦵ = − nFE⦵ is given in Section 1 of the data booklet and the value of 96 500 C mol-1 for Faraday's constant, F, is given in Section 2 of the data booklet. Students should understand what is meant by the term cells in series. International-mindednessMany of the world’s energy problems can be alleviated by using electrochemical cells as energy sources. Some cells, e.g. super-efficient microbial fuel cells, MFCs, (also known as biological fuel cells) can also help to clean-up the environment. The international community and national governments have to decide what to prioritise when it comes to funding research projects. |
Teaching tipsIt would be nice to be able to teach about standard redox potentials through relevant practical work - unfortunately I have yet to come across a school with its own standard hydrogen electrode. In fact relating all the E⦵ values to the hydrogen electrode seems strange to students (and perhaps teachers as well) as it is not immediately obvious why a gas (which does not conduct electricity) should be chosen as a standard. It is further complicated because most of the diagrams show hydrogen bubbling out from under the surface of the acid which means that the pressure must be greater than one atmosphere. What you can do though is build upon the core and construct an electrochemical series using Cu(s)/Cu2+(aq) as the standard and then shift all the values by 0.34 V after talking about why hydrogen is used. What is important is to impress upon students that the actual electrode potential is not known and cannot be measured - the Electrochemical Series is simply a table of relative values. Once they understand the concept they need plenty of practice at writing the relevant half-equations making sure that the half-cell with the more negative E⦵ value is the one that gives up electrons. They can also see that a species can act either as an oxidizing or reducing agent depending upon the standard electrode potential of the other half-cell. You will need to discuss the relationship between ΔG and E⦵ and bring in the Faraday constant, F. Explain that when the cell reaches equilibrium then both ΔG and E⦵ will be zero. Although students are given the value of F they do not need to use it to solve electrolysis calculations as the syllabus clearly says that they only need to determine relative amounts but they will need to realise that the same quantity of electricity passes through two different electrolytic cells when they are connected in series. The practical Electrolytic cells is well worth doing to reinforce the theory of electrolysis in solution. Students must realise that unlike a voltaic cell an electrolytic cell is a non-spontaneous process. Electricity needs to be put in to drive the reaction in the opposite direction to what happens in a voltaic cell. Too often when marking Paper 2 I see diagrams of electrolytic cells with voltmeters and salt-bridges and electrochemical cells powered by an external battery! | Study guide
Pages 74 - 76 QuestionsFor ten 'quiz' multiple choice questions with the answers explained see MC test: Electrochemical cells (AHL). For short-answer questions on voltaic cells which can be set as an assignment for a test, homework or given for self study together with model answers see Electrochemical cells (1) (AHL) questions and for a second set of short-answer questions on electrolytic cells see Electrochemical cells (2) (AHL) questions. Vocabulary list:Standard electrode potential, E IM, TOK, 'Utilization' etc.See separate page which covers all of Topic 9 & 19. Practical work |
Teaching slides
Teachers may wish to share these slides with students for learning or for reviewing key concepts.
Electrochemical cells (AHL):
Electrochemical cells (AHL):
Electrochemical cells (AHL):
Electrochemical cells (AHL):
Electrochemical cells (AHL):
Electrochemical cells (AHL):
Electrochemical cells (AHL):
Electrochemical cells (AHL):
Electrochemical cells (AHL):
Electrochemical cells (AHL):
Electrochemical cells (AHL):
Electrochemical cells (AHL):
Electrochemical cells (AHL):
Electrochemical cells (AHL):
Electrochemical cells (AHL):
Electrochemical cells (AHL):
Other resources
1. A useful video describing the standard hydrogen electrode. Note however that although written on the screen there is no verbal inclusion of the need for the temperature to be 25 oC (298 K) for standard conditions. Also the hydrogen gas shown must actually be slightly higher than the required one atmosphere pressure as it bubbles out from below the surface of the solution.
2. This video shows exactly how some people make the whole thing so unnecessarily complicated using cell conventions and left and right hands and negative EMF values that it is not surprising that students find the topic so impossibly difficult. Not recommended to show to students.
3. A good video on the electrolysis of concentrated sodium chloride solution which explains it well.
4. The electrolysis of 'water' using dilute sulfuric acid and a Hofmann voltameter. It clearly show the volume relationship and the identity of the two gases.