4.1.1 Electrolysis and electrolytic cells

What electrolysis is

Definition

Electrolysis

The breaking down of an ionic compound, when molten or in aqueous solution, by passing an electric current through it.

Electrolysis uses electrical energy to break an ionic compound down into simpler substances, which are usually elements.
It is a chemical change, because new substances are formed at the electrodes and the original compound is used up.
Electrolysis only works when the ionic compound is molten or dissolved in water to make an aqueous solution.
The current used is always a direct current (d.c.), so that one electrode stays positive and the other stays negative for the whole process.
An ionic compound can be broken down this way because it is made of ions, which are particles that carry an electric charge.
Electrolysis is used on a large scale to extract reactive metals such as aluminium, to manufacture chlorine and hydrogen, and to coat objects with a thin layer of metal.

Molten and aqueous electrolytes

Definition

Electrolyte

A molten or aqueous ionic compound that conducts electricity and is broken down during electrolysis.

The substance that is broken down during electrolysis is called the electrolyte.
An electrolyte conducts electricity because it contains ions that are free to move.
In a solid ionic compound, the ions are held in fixed positions in a giant ionic lattice, so they cannot move and the solid does not conduct.
When the compound is heated until it melts, the lattice breaks down and the ions can slide past one another.
When the compound dissolves in water, the lattice also breaks apart and the ions spread out through the water, where they are free to move.
Solid sodium chloride therefore does not conduct electricity, but both molten sodium chloride and aqueous sodium chloride do.
A molten electrolyte must be kept above its melting point, and sodium chloride, for example, has to be heated to $∘C801\ ^{\circ}\text{C}$ before it melts.
An aqueous electrolyte usually needs no heating, which makes it easier and safer to use in the laboratory.
Substances made of covalent molecules, such as sugar, ethanol and pure water, contain no ions or almost none, so they are not electrolytes.
A metal conducts electricity through its delocalised electrons rather than through ions, so it carries a current without being broken down and is not an electrolyte.

The electrolytic cell

Definition

Electrolytic cell

An arrangement of two electrodes dipping into an electrolyte and connected to a direct current power supply, used to carry out electrolysis.

Definition

Electrode

A solid conductor through which an electric current enters or leaves an electrolyte.

Definition

Inert electrode

An electrode, usually made of graphite or platinum, that conducts electricity without reacting with the electrolyte or the products.

Electrolysis is carried out in an electrolytic cell, which is made of four main parts.
1. Power supply: a d.c. source, such as a battery or a laboratory power pack, that drives the current around the circuit.
2. Electrodes: two solid rods or plates of a conducting material that dip into the electrolyte.
3. Electrolyte: the molten or aqueous ionic compound that fills the space between the electrodes.
4. Connecting wires: these join each electrode to one terminal of the power supply.
The circuit is only complete when the current passes through the electrolyte from one electrode to the other.
The two electrodes must not touch, otherwise the current passes straight between them and the electrolyte is bypassed.
A lamp or an ammeter is often added to the circuit to show that a current is flowing.
Electrodes are usually made of graphite, a form of carbon, or of platinum, because both conduct electricity well.
Graphite and platinum are described as inert, because they do not react with the electrolyte or with the substances formed on them.
Graphite is cheap and has a very high melting point, so it is used in most school experiments and for molten electrolytes.
Platinum is extremely unreactive but expensive, so it is used mainly in small laboratory cells.
For a molten electrolyte, the solid compound is placed in a heat-resistant container such as a crucible and heated strongly until it melts around the electrodes.
For an aqueous electrolyte, the solution is held in a beaker or a small plastic cell, and any gases formed can be collected in test tubes placed over the electrodes.

Anode and cathode

Definition

Anode

The positive electrode in an electrolytic cell.

Definition

Cathode

The negative electrode in an electrolytic cell.

The two electrodes are named according to the terminal of the power supply that each one is connected to.
The anode is the positive electrode, because it is connected to the positive terminal of the power supply.
The cathode is the negative electrode, because it is connected to the negative terminal of the power supply.
In a circuit diagram, the longer line of the cell symbol is the positive terminal, so the electrode wired to it is the anode.
The electrode wired to the shorter line of the cell symbol is the cathode.
Swapping the two wires at the power supply swaps the anode and the cathode, so always trace the wires rather than relying on where an electrode sits in the diagram.
The memory aid PANIC, which stands for Positive Anode Negative Is Cathode, helps you keep the two names the right way round.

Practical

Testing substances to find electrolytes

Aim: to find out which substances conduct electricity and are broken down by it, and so identify which of them are electrolytes.
Apparatus: low-voltage d.c. power pack set to $V6\ \text{V}$ , two graphite electrodes in an electrode holder, connecting leads with crocodile clips, a small filament lamp, a $cm3100\ \text{cm}^3$ beaker, a wash bottle of distilled water, paper towels, eye protection, and samples of solid sodium chloride, sodium chloride solution, copper(II) sulfate solution, dilute hydrochloric acid, sugar solution, ethanol and distilled water.
Method:
- Connect the power pack, the lamp and the two graphite electrodes in series, so that the gap between the electrodes is the only break in the circuit.
- Touch the electrodes together briefly to check that the lamp lights, then separate them again.
- Pour about $cm350\ \text{cm}^3$ of the first liquid into the beaker and lower the electrodes until they dip at least $cm2\ \text{cm}$ into it without touching each other.
- Switch on and watch the lamp for about one minute, looking closely at both electrodes for bubbles, a coloured coating or a change of colour in the liquid.
- Switch off, lift out the electrodes, rinse them with distilled water and dry them on a paper towel so that nothing is carried over into the next sample.
- Repeat with each liquid in turn, and finally push the electrodes into a heap of dry sodium chloride crystals.
Results:
- With sodium chloride solution, copper(II) sulfate solution and dilute hydrochloric acid, the lamp lights and bubbles or a coating appear at the electrodes, so these are electrolytes.
- With solid sodium chloride, sugar solution, ethanol and distilled water, the lamp stays off and nothing forms at the electrodes, so these are not electrolytes.
Conclusion: a substance is an electrolyte only when it contains ions that are free to move, which is why sodium chloride conducts when dissolved but not as a dry solid.
Watch out:
- Electrodes that are not rinsed between samples carry ions into the next liquid and can make a non-electrolyte appear to conduct.
- A very small current may not light the lamp, so an ammeter gives a more sensitive check than a lamp alone.
Safety: wear eye protection, keep ethanol well away from any flame because it is highly flammable, switch off promptly because small amounts of toxic chlorine can form with chloride solutions, and wash your hands after handling copper(II) sulfate.

Common Mistake

Do not say that solid ionic compounds conduct electricity, because their ions are fixed in the lattice and cannot move.
Do not call a metal wire an electrolyte, because it conducts by delocalised electrons and is not broken down.
Do not mix up the electrode names: in electrolysis the anode is always positive and the cathode is always negative.
Do not say that the electrodes are broken down, because it is the electrolyte that is decomposed.

Exam technique

Defining electrolysis and labelling cells

A full definition of electrolysis needs three ideas: an ionic compound is broken down, it is molten or in aqueous solution, and an electric current passes through it.
Leaving out the word ionic, or forgetting to say that the compound is molten or dissolved, is the most common way to lose the definition mark.
Before labelling a cell diagram, trace each wire back to the power supply and only then write anode or cathode.
When asked for an inert electrode material, name graphite or platinum and give the reason that it does not react.

Self review

What is electrolysis?
Why does solid sodium chloride not conduct electricity while aqueous sodium chloride does?
Which electrode is the anode, and which terminal of the power supply is it connected to?
Name two materials used to make inert electrodes.
Why is a copper wire not described as an electrolyte, even though it conducts electricity?

4.1.2 Products and observations in electrolysis

Where the products form

During electrolysis, each ion in the electrolyte is attracted to the electrode with the opposite charge.
Positive ions move towards the negative cathode, and negative ions move towards the positive anode.
Metal ions and hydrogen ions are positive, so metals or hydrogen are always formed at the cathode.
Ions of non-metals, such as chloride, bromide and oxide ions, are negative, so non-metals other than hydrogen are always formed at the anode.
When the ions reach the electrodes they are changed into neutral atoms or molecules, and these are the products that you see.
Using inert electrodes made of platinum or graphite means the electrodes take no part in the reaction, so every product comes from the electrolyte.
In a molten compound, the only ions present are the ions of that compound.
In an aqueous solution, water also provides a small number of hydrogen ions, $H+\text{H}^{+}$ , and hydroxide ions, $OH−\text{OH}^{-}$ , so hydrogen or oxygen can be formed instead of the elements in the dissolved compound.

Molten lead(II) bromide

Lead(II) bromide, $PbBr2\text{PbBr}_2$ , is a white ionic solid made of lead(II) ions, $Pb2+\text{Pb}^{2+}$ , and bromide ions, $Br−\text{Br}^{-}$ .
The solid is placed in a crucible with two graphite electrodes and heated strongly until it melts at about $∘C373\ ^{\circ}\text{C}$ .
A lamp in the circuit lights only once the solid has melted, because only then are the ions free to move.
The positive $Pb2+\text{Pb}^{2+}$ ions move to the cathode, where lead is formed.
The lead is produced as a silvery-grey liquid, because the electrolyte is hotter than the melting point of lead, $∘C327\ ^{\circ}\text{C}$ , and it collects as a shiny bead beneath the cathode.
The negative $Br−\text{Br}^{-}$ ions move to the anode, where bromine is formed.
The bromine escapes from the anode as red-brown fumes, because the temperature is far above the boiling point of bromine, $∘C59\ ^{\circ}\text{C}$ .
Overall, the compound is split into its two elements: $PbBr2(l)→Pb(l)+Br2(g)\text{PbBr}_2(\text{l}) \rightarrow \text{Pb}(\text{l}) + \text{Br}_2(\text{g})$ .
This electrolysis is carried out in a fume cupboard, because bromine vapour is toxic and corrosive and lead compounds are also toxic.

Concentrated aqueous sodium chloride

Concentrated aqueous sodium chloride, also called brine, contains sodium ions, $Na+\text{Na}^{+}$ , and chloride ions, $Cl−\text{Cl}^{-}$ , together with hydrogen ions and hydroxide ions from the water.
At the cathode, hydrogen is formed rather than sodium, because sodium is far too reactive to exist in contact with water.
You see bubbles of a colourless gas at the cathode, and the gas pops with a lighted splint, which shows that it is hydrogen.
At the anode, chlorine is formed from the chloride ions.
You see bubbles of a pale yellow-green gas at the anode, and the gas bleaches damp litmus paper, which shows that it is chlorine.
Chlorine is slightly soluble in water, so at first less gas collects at the anode than at the cathode.
The sodium ions and hydroxide ions left behind form sodium hydroxide, so the solution around the cathode becomes alkaline.
Industry uses this process to make three useful products from salt: hydrogen, chlorine and sodium hydroxide.

Dilute sulfuric acid

Dilute sulfuric acid, $H2SO4(aq)\text{H}_2\text{SO}_4(\text{aq})$ , contains hydrogen ions, $H+\text{H}^{+}$ , and sulfate ions, $SO42−\text{SO}_4^{2-}$ , together with a small number of hydroxide ions from the water.
At the cathode, hydrogen is formed, seen as bubbles of a colourless gas that pops with a lighted splint.
At the anode, oxygen is formed, seen as bubbles of a colourless gas that relights a glowing splint.
The sulfate ions are not changed, and the oxygen comes from the water instead.
The volume of hydrogen collected is twice the volume of oxygen, so the two gases form in a volume ratio of $2 : 1$ .
This ratio arises because the hydrogen and oxygen both come from water, which contains two hydrogen atoms for every oxygen atom: $2H2O(l)→2H2(g)+O2(g)2\text{H}_2\text{O}(\text{l}) \rightarrow 2\text{H}_2(\text{g}) + \text{O}_2(\text{g})$ .
As water is used up, the remaining acid slowly becomes more concentrated.
Pure water contains too few ions to conduct well, so the acid is added to provide plenty of mobile ions.

Practical

Electrolysis of aqueous solutions

Aim: to identify the products at each electrode when concentrated aqueous sodium chloride and dilute sulfuric acid are electrolysed using inert electrodes.
Apparatus: a small electrolysis cell with two graphite electrodes fitted through its base, two small test tubes, a low-voltage d.c. power pack, connecting leads, wooden splints, a Bunsen burner, blue litmus paper, universal indicator, concentrated sodium chloride solution, dilute sulfuric acid and eye protection.
Method, sodium chloride:
- Pour concentrated sodium chloride solution into the cell until the electrodes are covered by about $cm3\ \text{cm}$ of liquid, then add a few drops of universal indicator.
- Fill two small test tubes to the brim with the same solution, close each with a thumb, turn it upside down and lower it over one electrode without letting any air in.
- Connect the electrodes to the power pack set at about $V6\ \text{V}$ d.c. and note which electrode is joined to the negative terminal, as this is the cathode.
- Switch on and watch both electrodes, recording the colour of any gas and any change in the colour of the indicator.
- Switch off when one tube is nearly full of gas.
- Lift the cathode tube out still upside down and hold a lighted splint at its mouth.
- Hold a piece of damp blue litmus paper at the mouth of the anode tube.
Method, sulfuric acid:
- Empty and rinse the cell, then repeat the same steps with dilute sulfuric acid, leaving out the universal indicator.
- Compare the volumes of gas in the two tubes before switching off.
- Test the anode gas with a glowing splint instead of litmus paper.
Results:
- With sodium chloride, the colourless cathode gas pops with a lighted splint and the indicator near the cathode turns blue-purple, showing hydrogen and an alkaline solution.
- With sodium chloride, the pale yellow-green anode gas bleaches the damp litmus paper and the indicator near the anode loses its colour, showing chlorine.
- With sulfuric acid, both gases are colourless, the cathode tube holds about twice as much gas, the cathode gas pops and the anode gas relights a glowing splint.
Watch out:
- Air left in the tubes dilutes the gases and can make the splint tests fail.
- Hydrogen is less dense than air, so test it with the tube held mouth downwards or it escapes before the splint arrives.
- Chlorine dissolves in the solution, so the anode tube fills more slowly than the cathode tube at first.
Safety: wear eye protection, keep the volume of chlorine small and work in a well-ventilated room because chlorine is toxic, and rinse off any sulfuric acid that touches your skin.

Predicting products of molten compounds

Definition

Binary compound

A compound made of only two different elements.

A binary compound contains just two elements, such as sodium chloride, $NaCl\text{NaCl}$ , or zinc bromide, $ZnBr2\text{ZnBr}_2$ .
A molten binary ionic compound contains only two kinds of ion: one metal ion and one non-metal ion.
The metal is therefore always formed at the cathode.
The non-metal is always formed at the anode.
The halogens are released as molecules made of two atoms, so chlorine is $Cl2\text{Cl}_2$ , bromine is $Br2\text{Br}_2$ and iodine is $I2\text{I}_2$ .
Always name the anode product as the element, so the product from a chloride is chlorine and never chloride.
Molten sodium chloride gives sodium at the cathode and chlorine at the anode.
Molten potassium iodide gives potassium at the cathode and iodine at the anode.
These predictions apply only to the molten state, because dissolving the same compound in water adds hydrogen ions and hydroxide ions that can change the products.

Example

Molten zinc bromide

Zinc bromide, $ZnBr2\text{ZnBr}_2$ , is a binary ionic compound made of zinc ions, $Zn2+\text{Zn}^{2+}$ , and bromide ions, $Br−\text{Br}^{-}$ .
When it is molten there is no water present, so these are the only ions in the electrolyte.
The positive zinc ions are attracted to the cathode, so zinc forms at the cathode.
The negative bromide ions are attracted to the anode, so bromine forms at the anode and is seen as red-brown fumes.

$ZnBr2→Zn+Br2\text{ZnBr}_2 \rightarrow \text{Zn} + \text{Br}_2$

Example

Molten and aqueous sodium chloride

Molten sodium chloride contains only sodium ions and chloride ions.
It gives sodium at the cathode and chlorine at the anode.
Concentrated aqueous sodium chloride also contains hydrogen ions and hydroxide ions from water.
It gives hydrogen at the cathode and chlorine at the anode, while the sodium ions stay in solution.
The anode product is the same in both cases, but the cathode product changes because water is present.

Common Mistake

Do not name the anode product as bromide or chloride, because the ions are turned into the elements bromine and chlorine.
Do not give sodium as a product of aqueous sodium chloride, because hydrogen forms at the cathode instead.
Do not mix up the splint tests: hydrogen pops with a lighted splint, while oxygen relights a glowing splint.
Do not give the volume ratio for dilute sulfuric acid the wrong way round, because there is twice as much hydrogen as oxygen.

Exam technique

Describing products and observations

An observation is something you can see, so describe colours and states, such as bubbles of colourless gas or red-brown fumes, rather than just naming the product.
When a question asks for products and observations, give both: name the substance, then describe what is seen.
Quote gas tests in full: hydrogen pops with a lighted splint, oxygen relights a glowing splint and chlorine bleaches damp litmus paper.
Check whether the electrolyte is molten or aqueous before predicting anything, because that decides whether water can supply hydrogen or oxygen.

Self review

Which products form at each electrode when molten lead(II) bromide is electrolysed?
What would you see at the anode during the electrolysis of concentrated aqueous sodium chloride?
Why is the volume of hydrogen from dilute sulfuric acid twice the volume of oxygen?
Which type of element is always formed at the cathode?
Which products form at each electrode when molten potassium iodide is electrolysed?

4.1.3 Electroplating

Why objects are electroplated

Definition

Electroplating

The use of electrolysis to coat a metal object with a thin layer of another metal.

Definition

Corrosion

The gradual destruction of a metal as it reacts with substances around it, such as oxygen and water.

Electroplating uses electrolysis to cover a metal object with a thin, even layer of a different metal.
The coating is often only a few thousandths of a millimetre thick, so very little of an expensive metal is needed.
Metal objects are electroplated for two main reasons.
1. To improve their appearance: a shiny coating of a metal such as silver, gold or chromium makes an object look attractive and more valuable.
2. To improve their resistance to corrosion: a coating of a metal that does not corrode easily stops air and water reaching the metal underneath.
Electroplating lets a cheap, strong metal such as steel form the body of an object while a thin layer of a costly or decorative metal forms its surface.
Common examples of electroplating are shown below.
1. Silver plating: cutlery and trophies made from cheaper alloys are coated with silver so that they shine like solid silver.
2. Chromium plating: steel taps, bicycle handlebars and car trims are coated with chromium, which is hard, shiny and resists corrosion.
3. Gold plating: jewellery and electrical connectors are coated with gold, which looks attractive and does not tarnish.
4. Tin plating: steel food cans are coated with tin, which stops the steel rusting and does not react with the food.
5. Nickel plating: steel objects are coated with nickel to protect them from corrosion, often as a base layer underneath chromium.

Setting up an electroplating cell

Electroplating is carried out in an electrolytic cell connected to a d.c. power supply.
The object to be plated is made the cathode, so it is connected to the negative terminal.
A piece of the plating metal is made the anode, so it is connected to the positive terminal.
The electrolyte is an aqueous solution of a soluble compound of the plating metal, so it contains positive ions of that metal.
1. To plate with copper, the anode is copper and the electrolyte is aqueous copper(II) sulfate.
2. To plate with nickel, the anode is nickel and the electrolyte is aqueous nickel(II) sulfate.
3. To plate with silver, the anode is silver and the electrolyte is a solution containing silver ions.
The object must be cleaned thoroughly before plating, because grease, dirt or a layer of oxide on its surface stops the coating sticking evenly.
The object and the anode are placed facing each other in the electrolyte without touching.

How the coating forms

When the current is switched on, the positive metal ions in the electrolyte are attracted to the negative object.
At the surface of the object, the metal ions are turned into metal atoms, which build up as a thin coating.
At the same time, the metal anode slowly dissolves, releasing new metal ions into the electrolyte.
The ions used up at the cathode are replaced by ions from the anode, so the concentration of the electrolyte stays about the same.
Metal is therefore transferred from the anode to the object, so the anode loses mass and the object gains mass.
A small current passed for a longer time gives a smooth, even coating that sticks firmly.
A large current deposits the metal too quickly, giving a rough or powdery coating that flakes off easily.
The object is often turned during plating so that every side faces the anode and receives an even coating.
A coating only protects against corrosion while it stays unbroken, because a scratch lets air and water reach the metal underneath.

Practical

Copper plating a key

Aim: to coat a brass key with copper by electrolysis and measure the mass of copper deposited.
Apparatus: a small brass key, a strip of copper foil, a $cm3250\ \text{cm}^3$ beaker, about $cm3150\ \text{cm}^3$ of aqueous copper(II) sulfate, a low-voltage d.c. power pack, an ammeter, a variable resistor, connecting leads with crocodile clips, emery paper, washing-up liquid, distilled water, paper towels, forceps, a balance reading to $g0.01\ \text{g}$ and eye protection.
Method:
- Rub the key and the copper strip with emery paper until both are bright, which removes any surface oxide.
- Wash the key in warm water with a little washing-up liquid to remove grease, rinse it with distilled water and from then on handle it only with forceps.
- Dry the key and weigh it, recording the mass to $g0.01\ \text{g}$ .
- Pour the copper(II) sulfate solution into the beaker and hang the key and the copper strip on opposite sides, facing each other and not touching.
- Connect the key to the negative terminal and the copper strip to the positive terminal, with the ammeter and the variable resistor in series.
- Switch on and adjust the variable resistor until the current is about $A0.2\ \text{A}$ , then leave the cell running for $20$ minutes while keeping the current steady.
- Switch off, lift the key out, rinse it gently with distilled water and leave it to dry without rubbing its surface.
- Reweigh the key and subtract the starting mass to find the mass of copper deposited.
Results:
- The key becomes covered in a pink-brown layer of copper and its mass increases.
- The copper strip becomes thinner and loses mass.
- The blue colour of the solution stays the same, because the copper ions removed are replaced from the anode.
Watch out:
- A greasy or tarnished key gives a patchy coating that rubs off easily.
- A current much larger than about $A0.5\ \text{A}$ gives a dark, powdery deposit that does not stick.
- Wiping the key dry can remove some of the coating and make the mass gain too small.
Safety: wear eye protection, avoid skin contact with copper(II) sulfate solution because it is harmful and an irritant, and keep the voltage low.

Example

Silver plating a spoon

A spoon made from a cheap nickel alloy is to be given a silver surface.
The spoon is cleaned to remove grease and is then connected to the negative terminal, making it the cathode.
A bar of pure silver is connected to the positive terminal, making it the anode.
The electrolyte is a solution containing silver ions, $Ag+\text{Ag}^{+}$ .
Silver ions move to the spoon and form a thin, shiny layer of silver on its surface.
The silver anode dissolves and replaces the silver ions used, so the same electrolyte can be used again and again.
The finished spoon looks like solid silver for a small fraction of the cost.

Common Mistake

Do not connect the object to the positive terminal, because the coating only forms on the cathode.
Do not choose an electrolyte that lacks ions of the plating metal, such as sodium chloride solution for copper plating.
Do not say that electroplating makes an object stronger, because its purposes are to improve appearance and resistance to corrosion.

Exam technique

Describing electroplating

Describe electroplating with three points: the object is the cathode, the plating metal is the anode, and the electrolyte contains ions of the plating metal.
When a question names a metal, use that metal throughout your answer, for example a nickel anode and aqueous nickel(II) sulfate.
When asked why objects are electroplated, give both reasons: to improve appearance and to improve resistance to corrosion.

Self review

Give two reasons why metal objects are electroplated.
Which electrode is the object to be plated, and which terminal is it connected to?
What must the electrolyte contain when an object is plated with nickel?
Why does the anode lose mass during copper plating?
Why does a small current give a better coating than a large current?

4.1.4 Charge transfer, aqueous electrolysis and half-equations

How charge moves in electrolysis

Definition

Oxidation

The loss of electrons by a substance.

Definition

Reduction

The gain of electrons by a substance.

In an electrolytic cell, charge is carried by electrons in the external circuit and by ions in the electrolyte.
Movement of electrons in the external circuit:
1. The power supply pushes electrons out of its negative terminal, along the wire and onto the cathode.
2. At the same time, electrons are drawn off the anode and flow along the wire back to the positive terminal.
3. Electrons travel only through the wires and the electrodes, and never through the electrolyte itself.
Loss or gain of electrons at the electrodes:
1. At the cathode, positive ions gain electrons from the electrode, so reduction takes place.
2. At the anode, negative ions, or sometimes the metal of the anode itself, lose electrons to the electrode, so oxidation takes place.
3. Once the ions have gained or lost electrons they become neutral atoms or molecules, which is described as being discharged.
Movement of ions in the electrolyte:
1. Positive ions, or cations, move through the electrolyte towards the cathode.
2. Negative ions, or anions, move through the electrolyte towards the anode.
3. This movement of ions in opposite directions carries the charge between the two electrodes and completes the circuit.
For every electron gained at the cathode, one electron is released at the anode, so the same amount of charge passes through every part of the circuit.
The phrase OIL RIG, meaning Oxidation Is Loss and Reduction Is Gain of electrons, helps you link each electrode to the correct process.

Copper(II) sulfate with graphite electrodes

Aqueous copper(II) sulfate contains copper(II) ions, $Cu2+\text{Cu}^{2+}$ , and sulfate ions, $SO42−\text{SO}_4^{2-}$ , together with hydrogen ions, $H+\text{H}^{+}$ , and hydroxide ions, $OH−\text{OH}^{-}$ , from water.
At the cathode, copper is less reactive than hydrogen, so copper(II) ions are discharged instead of hydrogen ions.
A pink-brown layer of copper forms on the cathode: $Cu2+(aq)+2e−→Cu(s)\text{Cu}^{2+}(\text{aq}) + 2\text{e}^{-} \rightarrow \text{Cu}(\text{s})$ .
At the anode, sulfate ions are not discharged, so hydroxide ions lose electrons and oxygen is released instead.
Bubbles of a colourless gas that relights a glowing splint form at the anode: $4OH−(aq)→O2(g)+2H2O(l)+4e−4\text{OH}^{-}(\text{aq}) \rightarrow \text{O}_2(\text{g}) + 2\text{H}_2\text{O}(\text{l}) + 4\text{e}^{-}$ .
The blue colour of the solution fades, because copper(II) ions are removed at the cathode and nothing replaces them.
The solution slowly becomes acidic, because the hydrogen ions and sulfate ions left behind make sulfuric acid.
The graphite cathode gains mass as copper builds up on it, while the mass of the graphite anode does not change.

Copper(II) sulfate with copper electrodes

When both electrodes are made of copper, the cathode reaction is unchanged, and copper is deposited on the cathode.
At the anode, copper atoms lose electrons more readily than hydroxide ions do, so the anode itself dissolves: $Cu(s)→Cu2+(aq)+2e−\text{Cu}(\text{s}) \rightarrow \text{Cu}^{2+}(\text{aq}) + 2\text{e}^{-}$ .
The copper anode loses mass and becomes thinner, while the copper cathode gains mass.
No gas is given off at either electrode.
The blue colour stays the same, because every copper(II) ion removed at the cathode is replaced by one released at the anode.
When the anode is pure copper, the mass lost by the anode equals the mass gained by the cathode.
The same arrangement is used to purify copper, with an impure copper anode dissolving and pure copper building up on the cathode.

Practical

Electrolysis of copper(II) sulfate

Aim: to compare the products and the changes in electrode mass when aqueous copper(II) sulfate is electrolysed with graphite electrodes and with copper electrodes.
Apparatus: a $cm3250\ \text{cm}^3$ beaker, aqueous copper(II) sulfate of concentration about $mol/dm30.5\ \text{mol/dm}^3$ , two graphite rods, two strips of copper foil, a low-voltage d.c. power pack, an ammeter, a variable resistor, connecting leads with crocodile clips, emery paper, distilled water, paper towels, a balance reading to $g0.01\ \text{g}$ , a small test tube, wooden splints and eye protection.
Method, graphite electrodes:
- Weigh both graphite rods, then clamp them in the beaker of copper(II) sulfate solution about $cm3\ \text{cm}$ apart.
- Fill a small test tube with the solution and invert it over the anode to collect any gas.
- Connect the rods to the power pack through the ammeter and the variable resistor, switch on and set the current to about $A0.3\ \text{A}$ .
- Leave the cell running for $20$ minutes, keeping the current constant and watching both electrodes and the colour of the solution.
- Test the gas collected at the anode with a glowing splint.
- Switch off, rinse both rods gently with distilled water, let them dry and reweigh them.
Method, copper electrodes:
- Clean two copper strips with emery paper, rinse and dry them, then weigh each one and label which will be the anode and which the cathode.
- Repeat the electrolysis in fresh copper(II) sulfate solution using the same current for the same time.
- Rinse, dry and reweigh both copper strips.
Results:
- With graphite electrodes, pink-brown copper coats the cathode, a colourless gas that relights a glowing splint forms at the anode, and the blue colour fades.
- With copper electrodes, the cathode gains mass, the anode loses about the same mass, no gas forms and the blue colour stays the same.
Watch out:
- Keep the current and the time the same for both runs, because the mass of copper deposited depends on the total charge passed.
- Copper deposited too quickly is loose and can fall off during rinsing, which makes the gain in mass of the cathode look too small.
- Dab the electrodes dry or leave them to dry rather than rubbing them, so that no copper is wiped away.
Safety: wear eye protection, avoid skin contact with copper(II) sulfate solution because it is harmful and an irritant, and wash your hands at the end.

Electrolysis of aqueous halides

An aqueous halide, such as a chloride, bromide or iodide solution, contains the ions of the dissolved compound together with hydrogen ions and hydroxide ions from water.
There are two possible positive ions and two possible negative ions, so a separate decision is made at each electrode.
At the cathode:
1. If the metal is more reactive than hydrogen, such as sodium, potassium or magnesium, the metal ions stay in solution and hydrogen is produced.
2. If the metal is less reactive than hydrogen, such as copper or silver, the metal itself is produced.
3. The concentration of the solution has no effect on which product forms at the cathode.
At the anode:
1. In a concentrated halide solution, the halide ions are discharged and the halogen is produced, so a chloride gives chlorine, a bromide gives bromine and an iodide gives iodine.
2. In a dilute halide solution, hydroxide ions are discharged instead and oxygen is produced.
3. Chlorine appears as a pale yellow-green gas, bromine colours the solution around the anode orange-brown, and iodine colours it brown.
Concentrated aqueous sodium chloride therefore gives hydrogen and chlorine, while dilute aqueous sodium chloride gives hydrogen and oxygen.

Example

Concentrated aqueous potassium bromide

The solution contains potassium ions, bromide ions, hydrogen ions and hydroxide ions.
Potassium is more reactive than hydrogen, so hydrogen is produced at the cathode.

2\text{H}^{+}(\text{aq}) + 2\text{e}^{-} \rightarrow \text{H}_2(\text{g})

The solution is concentrated, so bromide ions are discharged and bromine is produced at the anode.

2\text{Br}^{-}(\text{aq}) \rightarrow \text{Br}_2(\text{aq}) + 2\text{e}^{-}

You see colourless bubbles at the cathode and an orange-brown colour spreading out from the anode.

Example

Dilute aqueous copper(II) chloride

The solution contains copper(II) ions, chloride ions, hydrogen ions and hydroxide ions.
Copper is less reactive than hydrogen, so copper is produced at the cathode.

\text{Cu}^{2+}(\text{aq}) + 2\text{e}^{-} \rightarrow \text{Cu}(\text{s})

The solution is dilute, so hydroxide ions are discharged and oxygen is produced at the anode.

4\text{OH}^{-}(\text{aq}) \rightarrow \text{O}_2(\text{g}) + 2\text{H}_2\text{O}(\text{l}) + 4\text{e}^{-}

You see a pink-brown coating on the cathode and colourless bubbles at the anode that relight a glowing splint.

Writing ionic half-equations

Definition

Ionic half-equation

An equation that shows the electrons gained or lost by one type of particle at one electrode.

An ionic half-equation shows what happens to one type of particle at one electrode, with the electrons written in.
Build each half-equation in four steps.
1. Write the formula of the particle that reacts and of the product, remembering that hydrogen, oxygen and the halogens form diatomic molecules.
2. Balance the atoms of each element by placing numbers in front of the formulae.
3. Balance the charge by adding electrons, $e−\text{e}^{-}$ , to the side that is more positive.
4. Check that the total charge is the same on both sides, and add state symbols if the question asks for them.
At the cathode, the electrons appear on the left-hand side because they are gained, so every cathode half-equation shows reduction.
At the anode, the electrons appear on the right-hand side because they are lost, so every anode half-equation shows oxidation.
The half-equations you need most often are listed below.
1. Lead at the cathode: $Pb2++2e−→Pb\text{Pb}^{2+} + 2\text{e}^{-} \rightarrow \text{Pb}$
2. Copper at the cathode: $Cu2++2e−→Cu\text{Cu}^{2+} + 2\text{e}^{-} \rightarrow \text{Cu}$
3. Hydrogen at the cathode: $2H++2e−→H22\text{H}^{+} + 2\text{e}^{-} \rightarrow \text{H}_2$
4. Sodium at the cathode from a molten compound: $Na++e−→Na\text{Na}^{+} + \text{e}^{-} \rightarrow \text{Na}$
5. Chlorine at the anode: $2Cl−→Cl2+2e−2\text{Cl}^{-} \rightarrow \text{Cl}_2 + 2\text{e}^{-}$
6. Bromine at the anode: $2Br−→Br2+2e−2\text{Br}^{-} \rightarrow \text{Br}_2 + 2\text{e}^{-}$
7. Oxygen at the anode: $4OH−→O2+2H2O+4e−4\text{OH}^{-} \rightarrow \text{O}_2 + 2\text{H}_2\text{O} + 4\text{e}^{-}$
8. Copper anode dissolving: $Cu→Cu2++2e−\text{Cu} \rightarrow \text{Cu}^{2+} + 2\text{e}^{-}$

Example

Building the oxygen half-equation

Hydroxide ions, $OH−\text{OH}^{-}$ , are discharged to form oxygen, $O2\text{O}_2$ , and water, $H2O\text{H}_2\text{O}$ .

\text{OH}^{-} \rightarrow \text{O}_2 + \text{H}_2\text{O}

One $O2\text{O}_2$ molecule and two $H2O\text{H}_2\text{O}$ molecules contain four oxygen atoms and four hydrogen atoms, so four hydroxide ions are needed.

4\text{OH}^{-} \rightarrow \text{O}_2 + 2\text{H}_2\text{O}

The left-hand side has a total charge of $4 -$ and the right-hand side has a charge of $0$ , so four electrons are added to the right.

4\text{OH}^{-} \rightarrow \text{O}_2 + 2\text{H}_2\text{O} + 4\text{e}^{-}

The electrons are on the right, so the half-equation shows oxidation, which matches a reaction at the anode.

Common Mistake

Do not say that electrons flow through the electrolyte, because ions carry the charge through the electrolyte and electrons only move in the wires and electrodes.
Do not put the electrons on the wrong side, because they are gained on the left at the cathode and lost on the right at the anode.
Do not use concentration to decide the cathode product, because only the reactivity of the metal compared with hydrogen decides it.
Do not write $Cl\text{Cl}$ or $H\text{H}$ as products, because chlorine and hydrogen are released as diatomic molecules, $Cl2\text{Cl}_2$ and $H2\text{H}_2$ .

Exam technique

Explaining charge transfer

For a charge-transfer question, describe three things in order: electrons moving in the external circuit, electrons being lost or gained at each electrode, and ions moving through the electrolyte.
When explaining the colour of copper(II) sulfate solution, link it to whether the copper(II) ions are replaced, so it fades with graphite and stays blue with copper.
Check every half-equation for both atoms and charge before moving on, because a missing electron loses the mark.
When asked why a change is oxidation or reduction, refer to loss or gain of electrons rather than to oxygen.

Self review

How is charge carried through the wires and through the electrolyte during electrolysis?
Why does the blue colour of copper(II) sulfate fade with graphite electrodes but not with copper electrodes?
Which products form at each electrode when dilute aqueous sodium chloride is electrolysed?
What is the ionic half-equation for the formation of bromine at the anode?
At which electrode does reduction take place, and why?

4.1.1 Electrolysis and electrolytic cells

What electrolysis is

Definition

Electrolysis

The breaking down of an ionic compound, when molten or in aqueous solution, by passing an electric current through it.

Electrolysis uses electrical energy to break an ionic compound down into simpler substances, which are usually elements.
It is a chemical change, because new substances are formed at the electrodes and the original compound is used up.
Electrolysis only works when the ionic compound is molten or dissolved in water to make an aqueous solution.
The current used is always a direct current (d.c.), so that one electrode stays positive and the other stays negative for the whole process.
An ionic compound can be broken down this way because it is made of ions, which are particles that carry an electric charge.
Electrolysis is used on a large scale to extract reactive metals such as aluminium, to manufacture chlorine and hydrogen, and to coat objects with a thin layer of metal.

Molten and aqueous electrolytes

Definition

Electrolyte

A molten or aqueous ionic compound that conducts electricity and is broken down during electrolysis.

The substance that is broken down during electrolysis is called the electrolyte.
An electrolyte conducts electricity because it contains ions that are free to move.
In a solid ionic compound, the ions are held in fixed positions in a giant ionic lattice, so they cannot move and the solid does not conduct.
When the compound is heated until it melts, the lattice breaks down and the ions can slide past one another.
When the compound dissolves in water, the lattice also breaks apart and the ions spread out through the water, where they are free to move.
Solid sodium chloride therefore does not conduct electricity, but both molten sodium chloride and aqueous sodium chloride do.
A molten electrolyte must be kept above its melting point, and sodium chloride, for example, has to be heated to $∘C801\ ^{\circ}\text{C}$ before it melts.
An aqueous electrolyte usually needs no heating, which makes it easier and safer to use in the laboratory.
Substances made of covalent molecules, such as sugar, ethanol and pure water, contain no ions or almost none, so they are not electrolytes.
A metal conducts electricity through its delocalised electrons rather than through ions, so it carries a current without being broken down and is not an electrolyte.

The electrolytic cell

Definition

Electrolytic cell

An arrangement of two electrodes dipping into an electrolyte and connected to a direct current power supply, used to carry out electrolysis.

Definition

Electrode

A solid conductor through which an electric current enters or leaves an electrolyte.

Definition

Inert electrode

An electrode, usually made of graphite or platinum, that conducts electricity without reacting with the electrolyte or the products.

Electrolysis is carried out in an electrolytic cell, which is made of four main parts.
1. Power supply: a d.c. source, such as a battery or a laboratory power pack, that drives the current around the circuit.
2. Electrodes: two solid rods or plates of a conducting material that dip into the electrolyte.
3. Electrolyte: the molten or aqueous ionic compound that fills the space between the electrodes.
4. Connecting wires: these join each electrode to one terminal of the power supply.
The circuit is only complete when the current passes through the electrolyte from one electrode to the other.
The two electrodes must not touch, otherwise the current passes straight between them and the electrolyte is bypassed.
A lamp or an ammeter is often added to the circuit to show that a current is flowing.
Electrodes are usually made of graphite, a form of carbon, or of platinum, because both conduct electricity well.
Graphite and platinum are described as inert, because they do not react with the electrolyte or with the substances formed on them.
Graphite is cheap and has a very high melting point, so it is used in most school experiments and for molten electrolytes.
Platinum is extremely unreactive but expensive, so it is used mainly in small laboratory cells.
For a molten electrolyte, the solid compound is placed in a heat-resistant container such as a crucible and heated strongly until it melts around the electrodes.
For an aqueous electrolyte, the solution is held in a beaker or a small plastic cell, and any gases formed can be collected in test tubes placed over the electrodes.

Anode and cathode

Definition

Anode

The positive electrode in an electrolytic cell.

Definition

Cathode

The negative electrode in an electrolytic cell.

The two electrodes are named according to the terminal of the power supply that each one is connected to.
The anode is the positive electrode, because it is connected to the positive terminal of the power supply.
The cathode is the negative electrode, because it is connected to the negative terminal of the power supply.
In a circuit diagram, the longer line of the cell symbol is the positive terminal, so the electrode wired to it is the anode.
The electrode wired to the shorter line of the cell symbol is the cathode.
Swapping the two wires at the power supply swaps the anode and the cathode, so always trace the wires rather than relying on where an electrode sits in the diagram.
The memory aid PANIC, which stands for Positive Anode Negative Is Cathode, helps you keep the two names the right way round.

Practical

Testing substances to find electrolytes

Aim: to find out which substances conduct electricity and are broken down by it, and so identify which of them are electrolytes.
Apparatus: low-voltage d.c. power pack set to $V6\ \text{V}$ , two graphite electrodes in an electrode holder, connecting leads with crocodile clips, a small filament lamp, a $cm3100\ \text{cm}^3$ beaker, a wash bottle of distilled water, paper towels, eye protection, and samples of solid sodium chloride, sodium chloride solution, copper(II) sulfate solution, dilute hydrochloric acid, sugar solution, ethanol and distilled water.
Method:
- Connect the power pack, the lamp and the two graphite electrodes in series, so that the gap between the electrodes is the only break in the circuit.
- Touch the electrodes together briefly to check that the lamp lights, then separate them again.
- Pour about $cm350\ \text{cm}^3$ of the first liquid into the beaker and lower the electrodes until they dip at least $cm2\ \text{cm}$ into it without touching each other.
- Switch on and watch the lamp for about one minute, looking closely at both electrodes for bubbles, a coloured coating or a change of colour in the liquid.
- Switch off, lift out the electrodes, rinse them with distilled water and dry them on a paper towel so that nothing is carried over into the next sample.
- Repeat with each liquid in turn, and finally push the electrodes into a heap of dry sodium chloride crystals.
Results:
- With sodium chloride solution, copper(II) sulfate solution and dilute hydrochloric acid, the lamp lights and bubbles or a coating appear at the electrodes, so these are electrolytes.
- With solid sodium chloride, sugar solution, ethanol and distilled water, the lamp stays off and nothing forms at the electrodes, so these are not electrolytes.
Conclusion: a substance is an electrolyte only when it contains ions that are free to move, which is why sodium chloride conducts when dissolved but not as a dry solid.
Watch out:
- Electrodes that are not rinsed between samples carry ions into the next liquid and can make a non-electrolyte appear to conduct.
- A very small current may not light the lamp, so an ammeter gives a more sensitive check than a lamp alone.
Safety: wear eye protection, keep ethanol well away from any flame because it is highly flammable, switch off promptly because small amounts of toxic chlorine can form with chloride solutions, and wash your hands after handling copper(II) sulfate.

Common Mistake

Do not say that solid ionic compounds conduct electricity, because their ions are fixed in the lattice and cannot move.
Do not call a metal wire an electrolyte, because it conducts by delocalised electrons and is not broken down.
Do not mix up the electrode names: in electrolysis the anode is always positive and the cathode is always negative.
Do not say that the electrodes are broken down, because it is the electrolyte that is decomposed.

Exam technique

Defining electrolysis and labelling cells

A full definition of electrolysis needs three ideas: an ionic compound is broken down, it is molten or in aqueous solution, and an electric current passes through it.
Leaving out the word ionic, or forgetting to say that the compound is molten or dissolved, is the most common way to lose the definition mark.
Before labelling a cell diagram, trace each wire back to the power supply and only then write anode or cathode.
When asked for an inert electrode material, name graphite or platinum and give the reason that it does not react.

Self review

What is electrolysis?
Why does solid sodium chloride not conduct electricity while aqueous sodium chloride does?
Which electrode is the anode, and which terminal of the power supply is it connected to?
Name two materials used to make inert electrodes.
Why is a copper wire not described as an electrolyte, even though it conducts electricity?

4.1.2 Products and observations in electrolysis

Where the products form

During electrolysis, each ion in the electrolyte is attracted to the electrode with the opposite charge.
Positive ions move towards the negative cathode, and negative ions move towards the positive anode.
Metal ions and hydrogen ions are positive, so metals or hydrogen are always formed at the cathode.
Ions of non-metals, such as chloride, bromide and oxide ions, are negative, so non-metals other than hydrogen are always formed at the anode.
When the ions reach the electrodes they are changed into neutral atoms or molecules, and these are the products that you see.
Using inert electrodes made of platinum or graphite means the electrodes take no part in the reaction, so every product comes from the electrolyte.
In a molten compound, the only ions present are the ions of that compound.
In an aqueous solution, water also provides a small number of hydrogen ions, $H+\text{H}^{+}$ , and hydroxide ions, $OH−\text{OH}^{-}$ , so hydrogen or oxygen can be formed instead of the elements in the dissolved compound.

Molten lead(II) bromide

Lead(II) bromide, $PbBr2\text{PbBr}_2$ , is a white ionic solid made of lead(II) ions, $Pb2+\text{Pb}^{2+}$ , and bromide ions, $Br−\text{Br}^{-}$ .
The solid is placed in a crucible with two graphite electrodes and heated strongly until it melts at about $∘C373\ ^{\circ}\text{C}$ .
A lamp in the circuit lights only once the solid has melted, because only then are the ions free to move.
The positive $Pb2+\text{Pb}^{2+}$ ions move to the cathode, where lead is formed.
The lead is produced as a silvery-grey liquid, because the electrolyte is hotter than the melting point of lead, $∘C327\ ^{\circ}\text{C}$ , and it collects as a shiny bead beneath the cathode.
The negative $Br−\text{Br}^{-}$ ions move to the anode, where bromine is formed.
The bromine escapes from the anode as red-brown fumes, because the temperature is far above the boiling point of bromine, $∘C59\ ^{\circ}\text{C}$ .
Overall, the compound is split into its two elements: $PbBr2(l)→Pb(l)+Br2(g)\text{PbBr}_2(\text{l}) \rightarrow \text{Pb}(\text{l}) + \text{Br}_2(\text{g})$ .
This electrolysis is carried out in a fume cupboard, because bromine vapour is toxic and corrosive and lead compounds are also toxic.

Concentrated aqueous sodium chloride

Concentrated aqueous sodium chloride, also called brine, contains sodium ions, $Na+\text{Na}^{+}$ , and chloride ions, $Cl−\text{Cl}^{-}$ , together with hydrogen ions and hydroxide ions from the water.
At the cathode, hydrogen is formed rather than sodium, because sodium is far too reactive to exist in contact with water.
You see bubbles of a colourless gas at the cathode, and the gas pops with a lighted splint, which shows that it is hydrogen.
At the anode, chlorine is formed from the chloride ions.
You see bubbles of a pale yellow-green gas at the anode, and the gas bleaches damp litmus paper, which shows that it is chlorine.
Chlorine is slightly soluble in water, so at first less gas collects at the anode than at the cathode.
The sodium ions and hydroxide ions left behind form sodium hydroxide, so the solution around the cathode becomes alkaline.
Industry uses this process to make three useful products from salt: hydrogen, chlorine and sodium hydroxide.

Dilute sulfuric acid

Dilute sulfuric acid, $H2SO4(aq)\text{H}_2\text{SO}_4(\text{aq})$ , contains hydrogen ions, $H+\text{H}^{+}$ , and sulfate ions, $SO42−\text{SO}_4^{2-}$ , together with a small number of hydroxide ions from the water.
At the cathode, hydrogen is formed, seen as bubbles of a colourless gas that pops with a lighted splint.
At the anode, oxygen is formed, seen as bubbles of a colourless gas that relights a glowing splint.
The sulfate ions are not changed, and the oxygen comes from the water instead.
The volume of hydrogen collected is twice the volume of oxygen, so the two gases form in a volume ratio of $2 : 1$ .
This ratio arises because the hydrogen and oxygen both come from water, which contains two hydrogen atoms for every oxygen atom: $2H2O(l)→2H2(g)+O2(g)2\text{H}_2\text{O}(\text{l}) \rightarrow 2\text{H}_2(\text{g}) + \text{O}_2(\text{g})$ .
As water is used up, the remaining acid slowly becomes more concentrated.
Pure water contains too few ions to conduct well, so the acid is added to provide plenty of mobile ions.

Practical

Electrolysis of aqueous solutions

Aim: to identify the products at each electrode when concentrated aqueous sodium chloride and dilute sulfuric acid are electrolysed using inert electrodes.
Apparatus: a small electrolysis cell with two graphite electrodes fitted through its base, two small test tubes, a low-voltage d.c. power pack, connecting leads, wooden splints, a Bunsen burner, blue litmus paper, universal indicator, concentrated sodium chloride solution, dilute sulfuric acid and eye protection.
Method, sodium chloride:
- Pour concentrated sodium chloride solution into the cell until the electrodes are covered by about $cm3\ \text{cm}$ of liquid, then add a few drops of universal indicator.
- Fill two small test tubes to the brim with the same solution, close each with a thumb, turn it upside down and lower it over one electrode without letting any air in.
- Connect the electrodes to the power pack set at about $V6\ \text{V}$ d.c. and note which electrode is joined to the negative terminal, as this is the cathode.
- Switch on and watch both electrodes, recording the colour of any gas and any change in the colour of the indicator.
- Switch off when one tube is nearly full of gas.
- Lift the cathode tube out still upside down and hold a lighted splint at its mouth.
- Hold a piece of damp blue litmus paper at the mouth of the anode tube.
Method, sulfuric acid:
- Empty and rinse the cell, then repeat the same steps with dilute sulfuric acid, leaving out the universal indicator.
- Compare the volumes of gas in the two tubes before switching off.
- Test the anode gas with a glowing splint instead of litmus paper.
Results:
- With sodium chloride, the colourless cathode gas pops with a lighted splint and the indicator near the cathode turns blue-purple, showing hydrogen and an alkaline solution.
- With sodium chloride, the pale yellow-green anode gas bleaches the damp litmus paper and the indicator near the anode loses its colour, showing chlorine.
- With sulfuric acid, both gases are colourless, the cathode tube holds about twice as much gas, the cathode gas pops and the anode gas relights a glowing splint.
Watch out:
- Air left in the tubes dilutes the gases and can make the splint tests fail.
- Hydrogen is less dense than air, so test it with the tube held mouth downwards or it escapes before the splint arrives.
- Chlorine dissolves in the solution, so the anode tube fills more slowly than the cathode tube at first.
Safety: wear eye protection, keep the volume of chlorine small and work in a well-ventilated room because chlorine is toxic, and rinse off any sulfuric acid that touches your skin.

Predicting products of molten compounds

Definition

Binary compound

A compound made of only two different elements.

A binary compound contains just two elements, such as sodium chloride, $NaCl\text{NaCl}$ , or zinc bromide, $ZnBr2\text{ZnBr}_2$ .
A molten binary ionic compound contains only two kinds of ion: one metal ion and one non-metal ion.
The metal is therefore always formed at the cathode.
The non-metal is always formed at the anode.
The halogens are released as molecules made of two atoms, so chlorine is $Cl2\text{Cl}_2$ , bromine is $Br2\text{Br}_2$ and iodine is $I2\text{I}_2$ .
Always name the anode product as the element, so the product from a chloride is chlorine and never chloride.
Molten sodium chloride gives sodium at the cathode and chlorine at the anode.
Molten potassium iodide gives potassium at the cathode and iodine at the anode.
These predictions apply only to the molten state, because dissolving the same compound in water adds hydrogen ions and hydroxide ions that can change the products.

Example

Molten zinc bromide

Zinc bromide, $ZnBr2\text{ZnBr}_2$ , is a binary ionic compound made of zinc ions, $Zn2+\text{Zn}^{2+}$ , and bromide ions, $Br−\text{Br}^{-}$ .
When it is molten there is no water present, so these are the only ions in the electrolyte.
The positive zinc ions are attracted to the cathode, so zinc forms at the cathode.
The negative bromide ions are attracted to the anode, so bromine forms at the anode and is seen as red-brown fumes.

$ZnBr2→Zn+Br2\text{ZnBr}_2 \rightarrow \text{Zn} + \text{Br}_2$

Example

Molten and aqueous sodium chloride

Molten sodium chloride contains only sodium ions and chloride ions.
It gives sodium at the cathode and chlorine at the anode.
Concentrated aqueous sodium chloride also contains hydrogen ions and hydroxide ions from water.
It gives hydrogen at the cathode and chlorine at the anode, while the sodium ions stay in solution.
The anode product is the same in both cases, but the cathode product changes because water is present.

Common Mistake

Do not name the anode product as bromide or chloride, because the ions are turned into the elements bromine and chlorine.
Do not give sodium as a product of aqueous sodium chloride, because hydrogen forms at the cathode instead.
Do not mix up the splint tests: hydrogen pops with a lighted splint, while oxygen relights a glowing splint.
Do not give the volume ratio for dilute sulfuric acid the wrong way round, because there is twice as much hydrogen as oxygen.

Exam technique

Describing products and observations

An observation is something you can see, so describe colours and states, such as bubbles of colourless gas or red-brown fumes, rather than just naming the product.
When a question asks for products and observations, give both: name the substance, then describe what is seen.
Quote gas tests in full: hydrogen pops with a lighted splint, oxygen relights a glowing splint and chlorine bleaches damp litmus paper.
Check whether the electrolyte is molten or aqueous before predicting anything, because that decides whether water can supply hydrogen or oxygen.

Self review

Which products form at each electrode when molten lead(II) bromide is electrolysed?
What would you see at the anode during the electrolysis of concentrated aqueous sodium chloride?
Why is the volume of hydrogen from dilute sulfuric acid twice the volume of oxygen?
Which type of element is always formed at the cathode?
Which products form at each electrode when molten potassium iodide is electrolysed?

4.1.3 Electroplating

Why objects are electroplated

Definition

Electroplating

The use of electrolysis to coat a metal object with a thin layer of another metal.

Definition

Corrosion

The gradual destruction of a metal as it reacts with substances around it, such as oxygen and water.

Electroplating uses electrolysis to cover a metal object with a thin, even layer of a different metal.
The coating is often only a few thousandths of a millimetre thick, so very little of an expensive metal is needed.
Metal objects are electroplated for two main reasons.
1. To improve their appearance: a shiny coating of a metal such as silver, gold or chromium makes an object look attractive and more valuable.
2. To improve their resistance to corrosion: a coating of a metal that does not corrode easily stops air and water reaching the metal underneath.
Electroplating lets a cheap, strong metal such as steel form the body of an object while a thin layer of a costly or decorative metal forms its surface.
Common examples of electroplating are shown below.
1. Silver plating: cutlery and trophies made from cheaper alloys are coated with silver so that they shine like solid silver.
2. Chromium plating: steel taps, bicycle handlebars and car trims are coated with chromium, which is hard, shiny and resists corrosion.
3. Gold plating: jewellery and electrical connectors are coated with gold, which looks attractive and does not tarnish.
4. Tin plating: steel food cans are coated with tin, which stops the steel rusting and does not react with the food.
5. Nickel plating: steel objects are coated with nickel to protect them from corrosion, often as a base layer underneath chromium.

Setting up an electroplating cell

Electroplating is carried out in an electrolytic cell connected to a d.c. power supply.
The object to be plated is made the cathode, so it is connected to the negative terminal.
A piece of the plating metal is made the anode, so it is connected to the positive terminal.
The electrolyte is an aqueous solution of a soluble compound of the plating metal, so it contains positive ions of that metal.
1. To plate with copper, the anode is copper and the electrolyte is aqueous copper(II) sulfate.
2. To plate with nickel, the anode is nickel and the electrolyte is aqueous nickel(II) sulfate.
3. To plate with silver, the anode is silver and the electrolyte is a solution containing silver ions.
The object must be cleaned thoroughly before plating, because grease, dirt or a layer of oxide on its surface stops the coating sticking evenly.
The object and the anode are placed facing each other in the electrolyte without touching.

How the coating forms

When the current is switched on, the positive metal ions in the electrolyte are attracted to the negative object.
At the surface of the object, the metal ions are turned into metal atoms, which build up as a thin coating.
At the same time, the metal anode slowly dissolves, releasing new metal ions into the electrolyte.
The ions used up at the cathode are replaced by ions from the anode, so the concentration of the electrolyte stays about the same.
Metal is therefore transferred from the anode to the object, so the anode loses mass and the object gains mass.
A small current passed for a longer time gives a smooth, even coating that sticks firmly.
A large current deposits the metal too quickly, giving a rough or powdery coating that flakes off easily.
The object is often turned during plating so that every side faces the anode and receives an even coating.
A coating only protects against corrosion while it stays unbroken, because a scratch lets air and water reach the metal underneath.

Practical

Copper plating a key

Aim: to coat a brass key with copper by electrolysis and measure the mass of copper deposited.
Apparatus: a small brass key, a strip of copper foil, a $cm3250\ \text{cm}^3$ beaker, about $cm3150\ \text{cm}^3$ of aqueous copper(II) sulfate, a low-voltage d.c. power pack, an ammeter, a variable resistor, connecting leads with crocodile clips, emery paper, washing-up liquid, distilled water, paper towels, forceps, a balance reading to $g0.01\ \text{g}$ and eye protection.
Method:
- Rub the key and the copper strip with emery paper until both are bright, which removes any surface oxide.
- Wash the key in warm water with a little washing-up liquid to remove grease, rinse it with distilled water and from then on handle it only with forceps.
- Dry the key and weigh it, recording the mass to $g0.01\ \text{g}$ .
- Pour the copper(II) sulfate solution into the beaker and hang the key and the copper strip on opposite sides, facing each other and not touching.
- Connect the key to the negative terminal and the copper strip to the positive terminal, with the ammeter and the variable resistor in series.
- Switch on and adjust the variable resistor until the current is about $A0.2\ \text{A}$ , then leave the cell running for $20$ minutes while keeping the current steady.
- Switch off, lift the key out, rinse it gently with distilled water and leave it to dry without rubbing its surface.
- Reweigh the key and subtract the starting mass to find the mass of copper deposited.
Results:
- The key becomes covered in a pink-brown layer of copper and its mass increases.
- The copper strip becomes thinner and loses mass.
- The blue colour of the solution stays the same, because the copper ions removed are replaced from the anode.
Watch out:
- A greasy or tarnished key gives a patchy coating that rubs off easily.
- A current much larger than about $A0.5\ \text{A}$ gives a dark, powdery deposit that does not stick.
- Wiping the key dry can remove some of the coating and make the mass gain too small.
Safety: wear eye protection, avoid skin contact with copper(II) sulfate solution because it is harmful and an irritant, and keep the voltage low.

Example

Silver plating a spoon

A spoon made from a cheap nickel alloy is to be given a silver surface.
The spoon is cleaned to remove grease and is then connected to the negative terminal, making it the cathode.
A bar of pure silver is connected to the positive terminal, making it the anode.
The electrolyte is a solution containing silver ions, $Ag+\text{Ag}^{+}$ .
Silver ions move to the spoon and form a thin, shiny layer of silver on its surface.
The silver anode dissolves and replaces the silver ions used, so the same electrolyte can be used again and again.
The finished spoon looks like solid silver for a small fraction of the cost.

Common Mistake

Do not connect the object to the positive terminal, because the coating only forms on the cathode.
Do not choose an electrolyte that lacks ions of the plating metal, such as sodium chloride solution for copper plating.
Do not say that electroplating makes an object stronger, because its purposes are to improve appearance and resistance to corrosion.

Exam technique

Describing electroplating

Describe electroplating with three points: the object is the cathode, the plating metal is the anode, and the electrolyte contains ions of the plating metal.
When a question names a metal, use that metal throughout your answer, for example a nickel anode and aqueous nickel(II) sulfate.
When asked why objects are electroplated, give both reasons: to improve appearance and to improve resistance to corrosion.

Self review

Give two reasons why metal objects are electroplated.
Which electrode is the object to be plated, and which terminal is it connected to?
What must the electrolyte contain when an object is plated with nickel?
Why does the anode lose mass during copper plating?
Why does a small current give a better coating than a large current?

4.1.4 Charge transfer, aqueous electrolysis and half-equations

How charge moves in electrolysis

Definition

Oxidation

The loss of electrons by a substance.

Definition

Reduction

The gain of electrons by a substance.

In an electrolytic cell, charge is carried by electrons in the external circuit and by ions in the electrolyte.
Movement of electrons in the external circuit:
1. The power supply pushes electrons out of its negative terminal, along the wire and onto the cathode.
2. At the same time, electrons are drawn off the anode and flow along the wire back to the positive terminal.
3. Electrons travel only through the wires and the electrodes, and never through the electrolyte itself.
Loss or gain of electrons at the electrodes:
1. At the cathode, positive ions gain electrons from the electrode, so reduction takes place.
2. At the anode, negative ions, or sometimes the metal of the anode itself, lose electrons to the electrode, so oxidation takes place.
3. Once the ions have gained or lost electrons they become neutral atoms or molecules, which is described as being discharged.
Movement of ions in the electrolyte:
1. Positive ions, or cations, move through the electrolyte towards the cathode.
2. Negative ions, or anions, move through the electrolyte towards the anode.
3. This movement of ions in opposite directions carries the charge between the two electrodes and completes the circuit.
For every electron gained at the cathode, one electron is released at the anode, so the same amount of charge passes through every part of the circuit.
The phrase OIL RIG, meaning Oxidation Is Loss and Reduction Is Gain of electrons, helps you link each electrode to the correct process.

Copper(II) sulfate with graphite electrodes

Aqueous copper(II) sulfate contains copper(II) ions, $Cu2+\text{Cu}^{2+}$ , and sulfate ions, $SO42−\text{SO}_4^{2-}$ , together with hydrogen ions, $H+\text{H}^{+}$ , and hydroxide ions, $OH−\text{OH}^{-}$ , from water.
At the cathode, copper is less reactive than hydrogen, so copper(II) ions are discharged instead of hydrogen ions.
A pink-brown layer of copper forms on the cathode: $Cu2+(aq)+2e−→Cu(s)\text{Cu}^{2+}(\text{aq}) + 2\text{e}^{-} \rightarrow \text{Cu}(\text{s})$ .
At the anode, sulfate ions are not discharged, so hydroxide ions lose electrons and oxygen is released instead.
Bubbles of a colourless gas that relights a glowing splint form at the anode: $4OH−(aq)→O2(g)+2H2O(l)+4e−4\text{OH}^{-}(\text{aq}) \rightarrow \text{O}_2(\text{g}) + 2\text{H}_2\text{O}(\text{l}) + 4\text{e}^{-}$ .
The blue colour of the solution fades, because copper(II) ions are removed at the cathode and nothing replaces them.
The solution slowly becomes acidic, because the hydrogen ions and sulfate ions left behind make sulfuric acid.
The graphite cathode gains mass as copper builds up on it, while the mass of the graphite anode does not change.

Copper(II) sulfate with copper electrodes

When both electrodes are made of copper, the cathode reaction is unchanged, and copper is deposited on the cathode.
At the anode, copper atoms lose electrons more readily than hydroxide ions do, so the anode itself dissolves: $Cu(s)→Cu2+(aq)+2e−\text{Cu}(\text{s}) \rightarrow \text{Cu}^{2+}(\text{aq}) + 2\text{e}^{-}$ .
The copper anode loses mass and becomes thinner, while the copper cathode gains mass.
No gas is given off at either electrode.
The blue colour stays the same, because every copper(II) ion removed at the cathode is replaced by one released at the anode.
When the anode is pure copper, the mass lost by the anode equals the mass gained by the cathode.
The same arrangement is used to purify copper, with an impure copper anode dissolving and pure copper building up on the cathode.

Practical

Electrolysis of copper(II) sulfate

Aim: to compare the products and the changes in electrode mass when aqueous copper(II) sulfate is electrolysed with graphite electrodes and with copper electrodes.
Apparatus: a $cm3250\ \text{cm}^3$ beaker, aqueous copper(II) sulfate of concentration about $mol/dm30.5\ \text{mol/dm}^3$ , two graphite rods, two strips of copper foil, a low-voltage d.c. power pack, an ammeter, a variable resistor, connecting leads with crocodile clips, emery paper, distilled water, paper towels, a balance reading to $g0.01\ \text{g}$ , a small test tube, wooden splints and eye protection.
Method, graphite electrodes:
- Weigh both graphite rods, then clamp them in the beaker of copper(II) sulfate solution about $cm3\ \text{cm}$ apart.
- Fill a small test tube with the solution and invert it over the anode to collect any gas.
- Connect the rods to the power pack through the ammeter and the variable resistor, switch on and set the current to about $A0.3\ \text{A}$ .
- Leave the cell running for $20$ minutes, keeping the current constant and watching both electrodes and the colour of the solution.
- Test the gas collected at the anode with a glowing splint.
- Switch off, rinse both rods gently with distilled water, let them dry and reweigh them.
Method, copper electrodes:
- Clean two copper strips with emery paper, rinse and dry them, then weigh each one and label which will be the anode and which the cathode.
- Repeat the electrolysis in fresh copper(II) sulfate solution using the same current for the same time.
- Rinse, dry and reweigh both copper strips.
Results:
- With graphite electrodes, pink-brown copper coats the cathode, a colourless gas that relights a glowing splint forms at the anode, and the blue colour fades.
- With copper electrodes, the cathode gains mass, the anode loses about the same mass, no gas forms and the blue colour stays the same.
Watch out:
- Keep the current and the time the same for both runs, because the mass of copper deposited depends on the total charge passed.
- Copper deposited too quickly is loose and can fall off during rinsing, which makes the gain in mass of the cathode look too small.
- Dab the electrodes dry or leave them to dry rather than rubbing them, so that no copper is wiped away.
Safety: wear eye protection, avoid skin contact with copper(II) sulfate solution because it is harmful and an irritant, and wash your hands at the end.

Electrolysis of aqueous halides

An aqueous halide, such as a chloride, bromide or iodide solution, contains the ions of the dissolved compound together with hydrogen ions and hydroxide ions from water.
There are two possible positive ions and two possible negative ions, so a separate decision is made at each electrode.
At the cathode:
1. If the metal is more reactive than hydrogen, such as sodium, potassium or magnesium, the metal ions stay in solution and hydrogen is produced.
2. If the metal is less reactive than hydrogen, such as copper or silver, the metal itself is produced.
3. The concentration of the solution has no effect on which product forms at the cathode.
At the anode:
1. In a concentrated halide solution, the halide ions are discharged and the halogen is produced, so a chloride gives chlorine, a bromide gives bromine and an iodide gives iodine.
2. In a dilute halide solution, hydroxide ions are discharged instead and oxygen is produced.
3. Chlorine appears as a pale yellow-green gas, bromine colours the solution around the anode orange-brown, and iodine colours it brown.
Concentrated aqueous sodium chloride therefore gives hydrogen and chlorine, while dilute aqueous sodium chloride gives hydrogen and oxygen.

Example

Concentrated aqueous potassium bromide

The solution contains potassium ions, bromide ions, hydrogen ions and hydroxide ions.
Potassium is more reactive than hydrogen, so hydrogen is produced at the cathode.

2\text{H}^{+}(\text{aq}) + 2\text{e}^{-} \rightarrow \text{H}_2(\text{g})

The solution is concentrated, so bromide ions are discharged and bromine is produced at the anode.

2\text{Br}^{-}(\text{aq}) \rightarrow \text{Br}_2(\text{aq}) + 2\text{e}^{-}

You see colourless bubbles at the cathode and an orange-brown colour spreading out from the anode.

Example

Dilute aqueous copper(II) chloride

The solution contains copper(II) ions, chloride ions, hydrogen ions and hydroxide ions.
Copper is less reactive than hydrogen, so copper is produced at the cathode.

\text{Cu}^{2+}(\text{aq}) + 2\text{e}^{-} \rightarrow \text{Cu}(\text{s})

The solution is dilute, so hydroxide ions are discharged and oxygen is produced at the anode.

4\text{OH}^{-}(\text{aq}) \rightarrow \text{O}_2(\text{g}) + 2\text{H}_2\text{O}(\text{l}) + 4\text{e}^{-}

You see a pink-brown coating on the cathode and colourless bubbles at the anode that relight a glowing splint.

Writing ionic half-equations

Definition

Ionic half-equation

An equation that shows the electrons gained or lost by one type of particle at one electrode.

An ionic half-equation shows what happens to one type of particle at one electrode, with the electrons written in.
Build each half-equation in four steps.
1. Write the formula of the particle that reacts and of the product, remembering that hydrogen, oxygen and the halogens form diatomic molecules.
2. Balance the atoms of each element by placing numbers in front of the formulae.
3. Balance the charge by adding electrons, $e−\text{e}^{-}$ , to the side that is more positive.
4. Check that the total charge is the same on both sides, and add state symbols if the question asks for them.
At the cathode, the electrons appear on the left-hand side because they are gained, so every cathode half-equation shows reduction.
At the anode, the electrons appear on the right-hand side because they are lost, so every anode half-equation shows oxidation.
The half-equations you need most often are listed below.
1. Lead at the cathode: $Pb2++2e−→Pb\text{Pb}^{2+} + 2\text{e}^{-} \rightarrow \text{Pb}$
2. Copper at the cathode: $Cu2++2e−→Cu\text{Cu}^{2+} + 2\text{e}^{-} \rightarrow \text{Cu}$
3. Hydrogen at the cathode: $2H++2e−→H22\text{H}^{+} + 2\text{e}^{-} \rightarrow \text{H}_2$
4. Sodium at the cathode from a molten compound: $Na++e−→Na\text{Na}^{+} + \text{e}^{-} \rightarrow \text{Na}$
5. Chlorine at the anode: $2Cl−→Cl2+2e−2\text{Cl}^{-} \rightarrow \text{Cl}_2 + 2\text{e}^{-}$
6. Bromine at the anode: $2Br−→Br2+2e−2\text{Br}^{-} \rightarrow \text{Br}_2 + 2\text{e}^{-}$
7. Oxygen at the anode: $4OH−→O2+2H2O+4e−4\text{OH}^{-} \rightarrow \text{O}_2 + 2\text{H}_2\text{O} + 4\text{e}^{-}$
8. Copper anode dissolving: $Cu→Cu2++2e−\text{Cu} \rightarrow \text{Cu}^{2+} + 2\text{e}^{-}$

Example

Building the oxygen half-equation

Hydroxide ions, $OH−\text{OH}^{-}$ , are discharged to form oxygen, $O2\text{O}_2$ , and water, $H2O\text{H}_2\text{O}$ .

\text{OH}^{-} \rightarrow \text{O}_2 + \text{H}_2\text{O}

One $O2\text{O}_2$ molecule and two $H2O\text{H}_2\text{O}$ molecules contain four oxygen atoms and four hydrogen atoms, so four hydroxide ions are needed.

4\text{OH}^{-} \rightarrow \text{O}_2 + 2\text{H}_2\text{O}

The left-hand side has a total charge of $4 -$ and the right-hand side has a charge of $0$ , so four electrons are added to the right.

4\text{OH}^{-} \rightarrow \text{O}_2 + 2\text{H}_2\text{O} + 4\text{e}^{-}

The electrons are on the right, so the half-equation shows oxidation, which matches a reaction at the anode.

Common Mistake

Do not say that electrons flow through the electrolyte, because ions carry the charge through the electrolyte and electrons only move in the wires and electrodes.
Do not put the electrons on the wrong side, because they are gained on the left at the cathode and lost on the right at the anode.
Do not use concentration to decide the cathode product, because only the reactivity of the metal compared with hydrogen decides it.
Do not write $Cl\text{Cl}$ or $H\text{H}$ as products, because chlorine and hydrogen are released as diatomic molecules, $Cl2\text{Cl}_2$ and $H2\text{H}_2$ .

Exam technique

Explaining charge transfer

For a charge-transfer question, describe three things in order: electrons moving in the external circuit, electrons being lost or gained at each electrode, and ions moving through the electrolyte.
When explaining the colour of copper(II) sulfate solution, link it to whether the copper(II) ions are replaced, so it fades with graphite and stays blue with copper.
Check every half-equation for both atoms and charge before moving on, because a missing electron loses the mark.
When asked why a change is oxidation or reduction, refer to loss or gain of electrons rather than to oxygen.

Self review

How is charge carried through the wires and through the electrolyte during electrolysis?
Why does the blue colour of copper(II) sulfate fade with graphite electrodes but not with copper electrodes?
Which products form at each electrode when dilute aqueous sodium chloride is electrolysed?
What is the ionic half-equation for the formation of bromine at the anode?
At which electrode does reduction take place, and why?

4.1 Electrolysis

4.1.1 Electrolysis and electrolytic cells

What electrolysis is

Molten and aqueous electrolytes

The electrolytic cell

Anode and cathode

Testing substances to find electrolytes

Defining electrolysis and labelling cells

4.1.2 Products and observations in electrolysis

Where the products form

Molten lead(II) bromide

Concentrated aqueous sodium chloride

Dilute sulfuric acid

Electrolysis of aqueous solutions

Predicting products of molten compounds

Molten zinc bromide

Molten and aqueous sodium chloride

Describing products and observations

4.1.3 Electroplating

Why objects are electroplated

Setting up an electroplating cell

How the coating forms

Copper plating a key

Silver plating a spoon

Describing electroplating

4.1.4 Charge transfer, aqueous electrolysis and half-equations

How charge moves in electrolysis

Copper(II) sulfate with graphite electrodes

Copper(II) sulfate with copper electrodes

Electrolysis of copper(II) sulfate

Electrolysis of aqueous halides

Concentrated aqueous potassium bromide

Dilute aqueous copper(II) chloride

Writing ionic half-equations

Building the oxygen half-equation

Explaining charge transfer

Review 4.1 Electrolysis by teaching Genie

Remember key concepts with flashcards

4.1 Electrolysis Revision Guide

4.1 Electrolysis

4.1.1 Electrolysis and electrolytic cells

What electrolysis is

Molten and aqueous electrolytes

The electrolytic cell

Anode and cathode

Testing substances to find electrolytes

Defining electrolysis and labelling cells

4.1.2 Products and observations in electrolysis

Where the products form

Molten lead(II) bromide

Concentrated aqueous sodium chloride

Dilute sulfuric acid

Electrolysis of aqueous solutions

Predicting products of molten compounds

Molten zinc bromide

Molten and aqueous sodium chloride

Describing products and observations

4.1.3 Electroplating

Why objects are electroplated

Setting up an electroplating cell

How the coating forms

Copper plating a key

Silver plating a spoon

Describing electroplating

4.1.4 Charge transfer, aqueous electrolysis and half-equations

How charge moves in electrolysis

Copper(II) sulfate with graphite electrodes

Copper(II) sulfate with copper electrodes

Electrolysis of copper(II) sulfate

Electrolysis of aqueous halides

Concentrated aqueous potassium bromide

Dilute aqueous copper(II) chloride

Writing ionic half-equations

Building the oxygen half-equation

Explaining charge transfer

Review 4.1 Electrolysis by teaching Genie

Remember key concepts with flashcards

4.1 Electrolysis Revision Guide