Syllabus coverage
- Define electrolysis as the decomposition of an ionic compound, when molten or in aqueous solution, by the passage of an electric current
- Identify in simple electrolytic cells the anode as the positive electrode, the cathode as the negative electrode, and the electrolyte as the molten or aqueous substance that undergoes electrolysis
- Identify the products formed at the electrodes and describe the observations made during the electrolysis of molten lead(II) bromide, concentrated aqueous sodium chloride and dilute sulfuric acid, using inert electrodes made of platinum or carbon (graphite)
- State that metals or hydrogen are formed at the cathode and that non-metals (other than hydrogen) are formed at the anode, and predict the identity of the products at each electrode for the electrolysis of a binary compound in the molten state
- Describe the transfer of charge during electrolysis, including the movement of electrons in the external circuit, the loss or gain of electrons at the electrodes, and the movement of ions in the electrolyte
- Identify the products formed at the electrodes and describe the observations made during the electrolysis of aqueous copper(II) sulfate using inert carbon (graphite) electrodes and when using copper electrodes
- Predict the identity of the products at each electrode for the electrolysis of a halide compound in dilute or concentrated aqueous solution
- Construct ionic half-equations for reactions at the anode (to show oxidation) and at the cathode (to show reduction)
- State that metal objects are electroplated to improve their appearance and resistance to corrosion, and describe how metals are electroplated
- State that a hydrogen-oxygen fuel cell uses hydrogen and oxygen to produce electricity with water as the only chemical product, and describe the advantages and disadvantages of using hydrogen-oxygen fuel cells in comparison with gasoline (petrol) engines in vehicles
Part of the IGCSE Co-ordinated Sciences syllabus 0654 / 0973.
What electrolysis is: electrolyte, electrodes and the circuit
Electrolysis is the decomposition of an ionic compound, when molten or in aqueous solution, by the passage of an electric current. The compound being broken down is the electrolyte, and the rods or plates that carry the current into it are the electrodes.
An ionic compound conducts electricity only when it is molten or dissolved in water. In the solid state the ions are held tightly in the lattice and cannot move, so there is nothing free to carry the charge. Melting or dissolving frees the ions, and it is this movement of ions that completes the circuit.
The two electrodes carry opposite charges. The anode is the positive electrode and the cathode is the negative electrode. Electrodes made of platinum or carbon (graphite) are described as inert: they carry the current but take no part in the reaction, so they are unchanged at the end and the electrolyte is the only substance decomposed.
- Electrolyte = the molten or aqueous ionic compound that undergoes electrolysis.
- Anode = the positive electrode; cathode = the negative electrode.
- The ions must be free to move, so a solid ionic compound will not conduct.
- Inert electrodes of platinum or carbon let you decompose the electrolyte without the electrode reacting too.
Products at the electrodes: what forms where
One simple rule predicts the products. Metals or hydrogen are formed at the cathode, because positive ions are attracted to the negative electrode. Non-metals other than hydrogen are formed at the anode, because negative ions are attracted to the positive electrode.
In a molten binary compound there is only one kind of positive ion and one kind of negative ion, so the prediction is straightforward. Molten lead(II) bromide, PbBr2, contains Pb2+ ions and Br- ions. The Pb2+ ions move to the cathode, the Br- ions move to the anode, and the compound is decomposed into its two elements.
At the cathode each lead ion gains two electrons and is reduced to lead metal. At the anode each bromide ion loses one electron and is oxidised to bromine. Two bromide ions are needed to make one bromine molecule, so the anode half-equation must be balanced with two electrons.
| Electrode | Ion and half-equation | Product and observation |
|---|---|---|
| Cathode (-) | Pb2+ + 2e- -> Pb | Silvery-grey beads of lead metal |
| Anode (+) | 2Br- -> Br2 + 2e- | Red-brown vapour of bromine |
| Overall | PbBr2 -> Pb + Br2 | The molten lead(II) bromide is decomposed into lead and bromine |
How the charge gets through: electrons, ions and the external circuit
Charge is carried in two different ways in two different places. In the external circuit - the wires and the power supply - charge is carried by electrons. In the electrolyte there are no free electrons, so charge is carried by ions.
The power supply pushes electrons into the cathode, making it negative, and pulls electrons out of the anode, making it positive. In the external circuit the electrons therefore flow from the anode, through the supply, to the cathode.
At the surface of each electrode electrons are transferred. At the cathode the positive ions collect electrons and are reduced; at the anode the negative ions give up their electrons and are oxidised. The electrons released at the anode travel round the external circuit and arrive at the cathode, where they are used again, so the circuit is complete.
Inside the electrolyte the positive ions (cations) move towards the cathode and the negative ions (anions) move towards the anode. This drift of ions is what carries the current through the melt or the solution.
- External circuit: electrons, flowing from the anode to the cathode through the power supply.
- At the electrodes: electrons are lost by ions at the anode (oxidation) and gained by ions at the cathode (reduction).
- Electrolyte: cations move to the cathode and anions move to the anode.
- The number of electrons lost at the anode equals the number gained at the cathode, so the charge balances overall.
Electrolysis of aqueous solutions: water joins in
When the electrolyte is dissolved in water, the water itself provides a small number of H+ and OH- ions. Those ions compete with the ions of the dissolved compound, so the products are not always what the formula of the compound alone would suggest.
Concentrated aqueous sodium chloride contains Na+, Cl-, H+ and OH- ions. At the cathode the H+ ions are reduced to hydrogen, because sodium is far too reactive to be deposited from solution. At the anode the chloride ions are plentiful, so they are oxidised to chlorine. Sodium hydroxide is left behind in the solution, so it becomes alkaline.
Dilute aqueous sodium chloride gives a different product at the anode. With only a few chloride ions present, the OH- ions from the water are oxidised instead, so oxygen is produced. Hydrogen is still formed at the cathode, and the solution again becomes alkaline.
Dilute sulfuric acid is the classic way to see the two gases together. Hydrogen is produced at the cathode and oxygen at the anode, and because the acid only supplies ions to carry the current, it is the water that is decomposed. The two gases are collected in a 2 : 1 volume ratio of hydrogen to oxygen, which matches the formula of water.
| Electrolyte | Cathode product | Anode product | What you see |
|---|---|---|---|
| Concentrated aqueous sodium chloride | Hydrogen | Chlorine | Colourless gas at the cathode that gives a squeaky pop with a lighted splint; pale green pungent gas at the anode that bleaches damp litmus paper |
| Dilute aqueous sodium chloride | Hydrogen | Oxygen | Colourless gas at the cathode; colourless gas at the anode that relights a glowing splint |
| Dilute sulfuric acid | Hydrogen | Oxygen | Both gases colourless, collected in a 2 : 1 volume ratio of hydrogen to oxygen |
Copper(II) sulfate: inert electrodes compared with copper electrodes
Copper(II) sulfate solution contains copper(II) ions, Cu2+, and sulfate ions. Which products you get depends entirely on what the electrodes are made of.
With inert carbon (graphite) electrodes, copper is deposited at the cathode and oxygen is produced at the anode. The cathode becomes coated with a pink-brown layer of copper, bubbles of colourless gas appear at the anode, and the blue colour of the solution slowly fades because the Cu2+ ions are being removed. The solution also becomes acidic, because sulfuric acid is left behind.
With copper electrodes the anode is no longer inert. Copper atoms at the anode lose electrons and dissolve as Cu2+ ions, so the anode loses mass. At the same time Cu2+ ions are reduced and deposited as copper on the cathode, so the cathode gains mass. The two changes cancel out, so the concentration of the copper(II) sulfate stays the same and the blue colour does not change.
That last arrangement is used to purify copper. Impure copper is made the anode and a thin sheet of pure copper is the cathode. Copper dissolves from the impure anode and is deposited as pure copper at the cathode, while the impurities fall to the bottom of the cell as a sludge.
| Electrode | Inert carbon electrodes | Copper electrodes |
|---|---|---|
| Cathode (-) | Cu2+ + 2e- -> Cu; pink-brown copper deposited | Cu2+ + 2e- -> Cu; copper deposited, so the cathode gains mass |
| Anode (+) | 4OH- -> O2 + 2H2O + 4e-; colourless oxygen gas given off | Cu -> Cu2+ + 2e-; the anode dissolves and loses mass |
| Solution | Blue colour fades and the solution becomes acidic | Concentration unchanged, so the blue colour stays the same |
Electroplating: using electrolysis to coat a metal
Electroplating means using electrolysis to put a thin layer of one metal onto the surface of another. Metal objects are electroplated to improve their appearance and to improve their resistance to corrosion.
The object to be plated is always made the cathode, the negative electrode, because positive metal ions are attracted to it. The anode is made of the plating metal, and the electrolyte is a solution of a salt of that same metal, so the ions being deposited are the ions of the plating metal.
When the current flows, metal ions from the solution are reduced and deposited as a thin layer on the object. The anode slowly dissolves and replaces the ions that have been used up, so the concentration of the electrolyte stays roughly constant and the plating keeps going.
Common examples are silver-plating cutlery and jewellery, chromium-plating taps and car parts for a hard shiny finish, nickel-plating to protect steel, and zinc-plating steel to make galvanised iron that resists rusting.
- Object to be plated = cathode (negative electrode).
- Plating metal = anode (positive electrode).
- Electrolyte = a solution of a salt of the plating metal.
- Purpose = improved appearance and better resistance to corrosion.
The hydrogen-oxygen fuel cell
A hydrogen-oxygen fuel cell uses hydrogen and oxygen to produce electricity, with water as the only chemical product. The overall reaction is the same as burning hydrogen, 2H2 + O2 -> 2H2O, but instead of releasing the energy as heat the cell converts it directly into electrical energy.
The great advantage over a petrol engine is what comes out of it. Burning petrol produces carbon dioxide, which contributes to global warming, as well as carbon monoxide and oxides of nitrogen. A hydrogen-oxygen fuel cell produces only water, so it does not pollute the air at the point of use.
There are real disadvantages too. Hydrogen is a gas that is highly flammable and forms an explosive mixture with air, so it must be stored and handled carefully. Storing it needs heavy, bulky high-pressure tanks and there are very few hydrogen refuelling stations. Fuel cells are also expensive because they need catalysts such as platinum, and most hydrogen is currently made from fossil fuels.
Compared with a petrol engine, a fuel cell is more efficient and has no moving parts, so it is quieter and needs less maintenance. Petrol, however, is a liquid that is easy to store and transport, has a high energy density, and there is already a filling station on every corner.
| Feature | Hydrogen-oxygen fuel cell | Petrol engine |
|---|---|---|
| Chemical product | Water only | Carbon dioxide and water, plus carbon monoxide and oxides of nitrogen |
| Efficiency | Higher, because energy is converted directly into electricity | Lower, because much of the energy is lost as heat |
| Pollution at the point of use | None | Polluting gases from the exhaust |
| Fuel storage | Hydrogen must be stored as a compressed gas or a liquid in heavy, bulky tanks | Petrol is a liquid, easy to store and transport |
| Refuelling | Very few hydrogen filling stations | A large existing network of filling stations |
| Moving parts | No moving parts, so quiet and low maintenance | Many moving parts, so noisier and needing more servicing |
Key terms you must be able to define
- Electrolysis
- The decomposition of an ionic compound, when molten or in aqueous solution, by the passage of an electric current.
- Electrolyte
- The molten or aqueous ionic compound that undergoes electrolysis; it carries the current by the movement of its ions.
- Electrode
- A conductor, such as a rod or plate of carbon or platinum, that carries the current into or out of the electrolyte.
- Anode
- The positive electrode, where negative ions lose electrons, so oxidation takes place.
- Cathode
- The negative electrode, where positive ions gain electrons, so reduction takes place.
- Oxidation
- The loss of electrons by a substance; during electrolysis it happens at the anode.
- Reduction
- The gain of electrons by a substance; during electrolysis it happens at the cathode.
- Electroplating
- Coating a metal object with a thin layer of another metal by making the object the cathode in an electrolytic cell, to improve its appearance and its resistance to corrosion.
Applied to the real world
Cambridge sets Co-ordinated Sciences in unfamiliar contexts - a hospital, a power station, a farm, a laboratory. Use these as models for how to apply the science, then write your own.
Concentrated brine is electrolysed on a huge scale. Chlorine from the anode is used to disinfect water supplies and to make bleach, PVC and solvents. Hydrogen from the cathode is used to make ammonia for fertiliser, and the sodium hydroxide left in the solution is used to make soap and to process paper.
Copper is purified by electrolysis using copper electrodes. Blocks of impure copper are the anodes and thin sheets of pure copper are the cathodes, with copper(II) sulfate solution as the electrolyte. Copper dissolves from the impure anode and is deposited as 99.99 per cent pure copper on the cathode, which is what electrical wiring needs.
Electroplating protects and decorates metal objects. Steel car bumpers are chromium-plated for a hard, shiny, corrosion-resistant finish, cheap cutlery is silver-plated, and steel is zinc-plated to make galvanised sheet that will not rust. In each case the object is the cathode and the plating metal is the anode.
Buses, cars and trains have been built with hydrogen-oxygen fuel cells. Hydrogen from a tank reacts with oxygen from the air to make electricity, and the only chemical product is water, so the exhaust is a trickle of pure water rather than carbon dioxide and other polluting gases.
Dilute sulfuric acid is electrolysed in a voltameter with inert platinum electrodes to split water into its elements. Hydrogen collects at the cathode and oxygen at the anode in a 2 : 1 volume ratio, which is a simple experimental confirmation that the formula of water is H2O.
Flashcards
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Exam-style questions
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Mark scheme — tick each point you actually made
Mark scheme — tick each point you actually made
Mark scheme — tick each point you actually made
Mark scheme — tick each point you actually made
Mark scheme — tick each point you actually made
Mark scheme — tick each point you actually made
Mark scheme — tick each point you actually made
Mark scheme — tick each point you actually made
Mark scheme — tick each point you actually made
Exam technique for this topic
- Learn the polarity first: in electrolysis the anode is the positive electrode and the cathode is the negative electrode. Reversing them costs marks in almost every part of an electrolysis question.
- Use OIL RIG and attach it to the electrodes: oxidation (loss of electrons) at the anode, reduction (gain of electrons) at the cathode.
- In a half-equation, the electrons go on the left for reduction at the cathode and on the right for oxidation at the anode. Check that both the atoms and the charges balance on each side.
- Whenever the electrolyte is an aqueous solution, decide what the water contributes. H+ ions are reduced to hydrogen at the cathode, and OH- ions are oxidised to oxygen at the anode when there are too few negative ions from the compound.
- Give the observation as well as the product: name the gas and give its test, or describe the colour of the deposit. 'A gas is given off' scores nothing.
- For fuel cell comparisons, structure your answer as advantages and disadvantages, and cover the exhaust gases, the efficiency, the storage of the fuel and the cost.
Why students lose marks here
Where this fits in IGCSE Co-ordinated Sciences
This page is one topic of Cambridge IGCSE Co-ordinated Sciences 0654 / 0973. IGCSE Co-ordinated Sciences is the course home: every area and every topic in Cambridge's own order, the papers you sit and what each one tests, and the exam-practice tools. The full topic list, with each topic linked to its revision page, is on the IGCSE Co-ordinated Sciences syllabus page.