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DIY: Electrolysis

Updated Thursday, 27th September 2007

A step-by-step guide to electrolysing water, part of the BBC/OU's programme website for Rough Science 1

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What is electrolysis?

In batteries, an electric current is created by chemical reactions that produce electrons (for more on batteries, visit 'How do batteries work?'). The reverse process, in which chemical reactions are produced by an electric current, is also possible. It's called electrolysis and it's something that we made frequent use of on our castaway island.

In electrolysis, a solution (the electrolyte) undergoes a chemical change when it conducts an electric current. The process is usually carried out in an apparatus consisting of connecting rods (called electrodes) dipped into the electrolyte. The beauty of electrolysis is that the energy in the electric current can be used to drive a whole range of so-called 'electrochemical' reactions that wouldn't normally occur spontaneously.

seaweed and jewellery The production of silver iodide for use in the castaway camera was one way in which we used electrolysis. Normally, silver prefers to exist in its metallic form, which is made up of neutral silver atoms. It’s pretty unreactive, as anyone with silver jewellery can testify. To make silver iodide we needed to produce positively charged silver ions from these silver atoms.

In other words, the silver atoms each had to lose an electron. Because of the unreactivity of silver metal, this wouldn’t happen very easily. We had to make it happen, using an electrolytic cell. We generated an electric current using an improvised seawater battery, and used the energy of that current to force the silver atoms in a pure silver bracelet electrode to give up their electrons. As this happened, we saw the silver disappear as the metal corroded away. The silver ions that formed conveniently dissolved in the electrolyte.

the seawater battery

Having produced silver ions, we then needed to react them with (negatively charged) iodide ions. The electrolyte that we used in our apparatus contained potassium iodide, which we’d extracted from seaweed. In aqueous solution, potassium iodide exists as separate potassium ions and iodide ions. So, as the current from the seawater battery generated the silver ions, they reacted spontaneously with the iodide ions already in the electrolyte, to give the compound we wanted: silver iodide. Clever, eh?

Electrolysis is used extensively in the extraction and purification of metals from their ores. Sodium metal, for instance, was first prepared in 1807, by the English chemist Sir Humphry Davy (who first coined the term electrochemistry). He did it by passing the strongest electrical current he could get (from three connected voltaic piles totalling 600 double plates) through molten sodium hydroxide (caustic soda). The metal potassium can be extracted in exactly the same way, using molten potassium hydroxide as the electrolyte.

Several industrially important metals, like sodium, magnesium and aluminium, are all produced commercially from the electrolysis of their molten salts.

But it’s not only the metallic products of electrolysis that are useful. For example, sodium metal and (industrially useful) chlorine gas are both produced when an electric current is passed through molten sodium chloride. Doing the same thing to an aqueous sodium chloride solution (seawater), on the other hand, produces sodium hydroxide as well as chlorine – another useful chemical.

In fact, electrolysis is used to manufacture these chemicals industrially. The energy in an electric current can even be used to break up water molecules into their constituent atoms: hydrogen and oxygen. The activity below shows you how.

Electrolysis can also be used to deposit one metal onto the surface of another in a process called electroplating. It’s generally done to improve the appearance or corrosion resistance of the base metal. The electrolyte used usually contains a high concentration of the metal ions to be plated. ‘Tin’ food cans are produced in this way. The ‘tin’ is actually thin steel plate that has been coated electrolytically with a very thin layer of tin to protect the steel from corroding.

Up to this point we have concentrated on some of the more useful aspects of electrochemistry; storing and obtaining electrical energy from batteries and producing useful chemicals by electrolysis. There is, however, a downside to the process: corrosion, which is responsible for millions of pounds worth of damage each year.

Nearly everywhere you see a corroded metal you are looking at the destructive result of electrochemical reactions. Many of the steps taken to prevent corrosion are aimed at breaking the electric circuit so as to stop these electrochemical reactions taking place.

When an electric current is passed through water in an electrolysis cell, tiny gas bubbles will form around each electrode. Hydrogen is given off at one electrode, and oxygen at the other.


  • a 2.5 litre plastic bottle
  • 2 x 0.5 litre plastic bottles (or smaller)
  • 2 pencils (the bigger the leads the better)
  • 1m of insulated copper wire
  • 500cm3 tap water (1cm3 is the same as 1ml)
  • 9 V battery
  • waterproof glue
  • splint or taper


  1. Cut the top off the large plastic bottle, to give a container about 15cm deep. Make two 10cm-deep collecting jars by doing the same thing to the smaller plastic bottles.
  2. Strip as much wood as you can from the pencils, so that you end up with two 10cm lengths of bare graphite. One way is to burn the wood over a candle. The graphite electrodes won’t corrode in tap water and interfere with the electrochemical reactions. Carefully make a small hole in the large container, and poke the electrodes through the holes. Make any gaps watertight using waterproof glue.
  3. Cut the copper wire in half and remove about 2cm of the insulating plastic from each of the four ends. Take one of the lengths of wire and wrap the bare copper at one end round the exposed end of one of the graphite electrodes. Make sure that there’s a good connection between the wire and the graphite.
  4. Repeat this with the other length of wire and electrode.
  5. Pour enough tap water into the large container so that the water level is about 4cm above the entry points of the electrodes. Invert a water-filled collecting jar over each of the electrodes, being careful to avoid losing any of the water.
  6. Now all you have to do is connect the exposed end of each length of wire to the terminals of a 9V battery. It doesn’t matter which way round you connect them. As the current passes through the cell, watch what happens – it may take some time before you see any results. If you live in a soft water area you may need to add a couple of drops of vinegar to the water to kick-start the electrolysis.
  7. The gas collected at one electrode is oxygen. You can test that this is the case by lighting a taper and then blowing it out and inserting it in the gas – the taper should relight. That produced at the other electrode will ignite with a pop, indicating that it’s hydrogen.


On Electrical Decomposition, by Michael Faraday
Fun Science Gallery - Galvanic Deposition
Batteries - Harnessing Redox Reactions
Science Toys You Can Make With Your Kids - Plastic Hydrogen Bomb


Block 6 of S103, Discovering Science, The Open University, 1998 ISBN 0 7492 8192 8

Snyder C.H., Chapter 6, The Extraordinary Chemistry of Everyday Things, 3rd edn., Wiley, 1998 ISBN 0 4711 7905 1

Faraday’s laws, corrosion and passivation’, in General Chemistry in the Laboratory, 3rd edn., Freeman (USA), 1991 ISBN 0 7167 2120 1


Here are some books and articles that you may want to try and get hold of:

Barrow J. D., The Artful Universe, Oxford University Press, 1995 ISBN 0 1985 3996 7.
A quite remarkable book that will change the way you view the world. Extremely accessible.

Burton et al., Chemical Storylines, G. Heinemann Educational Publishers, 1994 ISBN 0 435 63106 3.
Part of the Salters Advanced Chemistry course, which explores the frontiers of research and the applications of contemporary chemistry. For A level and other science courses aimed at 16 to 19-year olds.

Fraser A. and Gilchrist I., Starting Science (Book 1), Oxford University Press, 1998 ISBN 0 19 914235 1.
Part of an integrated science course for the National Curriculum Key Stage 3 and Scottish Environmental Studies (science) for S1 and S2.

Northedge A. et al., The Sciences Good Study Guide, The Open University, 1997 ISBN 0 7492 3411 3.
Indispensable for students of science, technology, mathematics and engineering. Packed with practical exercises and activities, all aimed at making studying more enjoyable and rewarding. Lots of hints and tips for those returning to study.

Selinger B., Chemistry in the Marketplace, 5th edn., Harcourt Brace, 1998 ISBN 0 7295 3300 X.
An excellent and informative reference source for all kinds of real-life applications of chemistry. Explores the world of chemistry that surrounds us in our daily lives, explained in terms that everyone can understand. ‘Makes chemistry come alive.’

PS547 Chemistry for Science Teachers course materials, The Open University, 1992
A course designed for use by science teachers from a wide variety of backgrounds, with varying experience of teaching science. A familiarity with some basic science (perhaps physics or biology) is assumed, but little understanding of chemistry is required. The mathematical understanding needed for the course is not great.





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