Electrolysis

Electrolysis uses electricity to break down a compound that could not be split apart easily any other way. It is how aluminium is pulled from its ore, how chlorine and sodium hydroxide are manufactured, and how metal objects are electroplated. This lesson, part of Topic C3 of OCR Gateway Combined Science, explains the apparatus and the key terms, works out the products of electrolysing molten and aqueous compounds, looks at the electrolysis of brine, and (for Higher tier) introduces the half-equations at each electrode.

By the end of this lesson you should be able to define electrolysis and its key terms, predict the products at each electrode for molten and aqueous compounds, describe the electrolysis of brine, and (Higher tier) write half-equations balanced for atoms and charge.

This lesson builds AO1 recall of the electrolysis key terms and apparatus, AO2 application when you predict the electrode products and (Higher) write balanced half-equations, and AO3 analysis when you work out which ion is discharged at each electrode for a given molten or aqueous compound.

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What Is Electrolysis?

Electrolysis is the breaking down of an ionic compound using an electric current. It only works when the ions are free to move — either because the compound is molten (melted) or dissolved in water — so that they can carry the charge and travel to the electrodes.

The key terms are:

• Electrolyte — the molten or dissolved ionic compound that conducts the current and is broken down.
• Electrodes — the two conducting rods (usually inert graphite or platinum) dipped into the electrolyte and connected to the power supply.
• Cathode — the negative electrode. Positive ions (cations) are attracted to it.
• Anode — the positive electrode. Negative ions (anions) are attracted to it.

Opposite charges attract, so positive ions travel to the negative cathode and negative ions travel to the positive anode. There they gain or lose electrons and are turned into neutral atoms or molecules — the products.

Exam Tip: Remember the polarity: cathode = negative, anode = positive. Positive ions go to the negative cathode; negative ions go to the positive anode. A handy memory aid is PANIC — Positive Anode, Negative Is Cathode.

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Electrolysis of Molten Ionic Compounds

When an ionic compound is molten, it contains only its own two ions, so the products are simple:

• the metal forms at the cathode (the positive metal ion gains electrons);
• the non-metal forms at the anode (the negative non-metal ion loses electrons).

Electrolysing molten lead bromide ($\text{PbBr}_2$PbBr2​), for example, gives lead at the cathode and bromine at the anode:

$\text{PbBr}_2 \rightarrow \text{Pb} + \text{Br}_2$PbBr2​→Pb+Br2​

This is exactly how aluminium is extracted industrially: molten aluminium oxide is electrolysed to give aluminium at the cathode and oxygen at the anode.

Exam Tip: For a molten compound the rule is simple — metal at the cathode, non-metal at the anode. There is no water to complicate matters, so the products are just the two elements from the compound.

Why electrolysis is used for reactive metals

Electrolysis links straight back to the reactivity series and to redox. A metal more reactive than carbon — such as aluminium — cannot be extracted by heating its oxide with carbon, because it grips its oxygen too tightly for carbon to remove. Instead, electrical energy is used to force the reduction: in electrolysis the metal ions gain electrons at the cathode (reduction) to form the metal.

For aluminium, the aluminium oxide is first melted so its ions are free to move. In industry it is also dissolved in molten cryolite to lower the melting point, which saves a great deal of energy (melting pure aluminium oxide would need a far higher temperature). The aluminium ions are then reduced to aluminium metal at the cathode. This is why extracting reactive metals by electrolysis is expensive — it uses a lot of electrical energy — and why recycling such metals is so worthwhile.

Exam Tip: Metals more reactive than carbon (e.g. aluminium) are extracted by electrolysis, not by reduction with carbon. Electrolysis is expensive because melting the compound and supplying the current both use a lot of energy — a strong argument for recycling.

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Electrolysis of Aqueous Solutions

When an ionic compound is dissolved in water, things get more complicated, because the water itself also supplies a few $\text{H}^+$H+ and $\text{OH}^-$OH− ions. Now there is a competition at each electrode, settled by these rules:

At the cathode (−):

• Hydrogen is produced unless the metal is less reactive than hydrogen (such as copper or silver), in which case the metal is produced.

At the anode (+):

• Oxygen is produced unless a halide ion (chloride, bromide or iodide) is present, in which case the halogen (chlorine, bromine or iodine) is produced.

Electrode	Rule	Copper sulfate solution	Sodium chloride solution
Cathode (−)	H₂ unless metal less reactive than H	Copper (less reactive than H)	Hydrogen (sodium more reactive than H)
Anode (+)	O₂ unless a halide is present	Oxygen (sulfate, no halide)	Chlorine (chloride is a halide)

So electrolysing copper sulfate solution with inert electrodes gives copper at the cathode and oxygen at the anode; electrolysing sodium chloride solution gives hydrogen at the cathode and chlorine at the anode.

Exam Tip: Apply the two aqueous rules in order. Cathode: is the metal more or less reactive than hydrogen? Less reactive (copper, silver) → the metal; more reactive (sodium, potassium, etc.) → hydrogen. Anode: is a halide present? Yes → the halogen; no → oxygen.

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Electrolysis of Brine

Brine is concentrated sodium chloride solution, and its electrolysis matters industrially because it produces three useful products:

• at the cathode (−): hydrogen (sodium is more reactive than hydrogen, so hydrogen is released);
• at the anode (+): chlorine (chloride is a halide, so the halogen is released);
• left in solution: sodium hydroxide (an alkali).

All three products are valuable: chlorine is used to sterilise water and to make bleach and plastics; hydrogen is used as a fuel and in making other chemicals; sodium hydroxide is used in soaps, paper-making and many industrial processes.

Exam Tip: Learn the three products of electrolysing brine: hydrogen (cathode), chlorine (anode) and sodium hydroxide (left in solution). A common misconception is to forget the sodium hydroxide because it is not formed at an electrode — but it is a key product.

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Higher Tier: Half-Equations

Higher tier: What happens at each electrode can be written as a half-equation, which shows the electrons ($e^-$e−) gained or lost. A half-equation must balance for atoms and charge.

• At the cathode, positive ions gain electrons (this is reduction). For example, hydrogen being produced:

$2\text{H}^+ + 2e^- \rightarrow \text{H}_2$2H++2e−→H2​

and copper being deposited:

$\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}$Cu2++2e−→Cu

• At the anode, negative ions lose electrons (this is oxidation). For example, chlorine being produced:

$2\text{Cl}^- \rightarrow \text{Cl}_2 + 2e^-$2Cl−→Cl2​+2e−

and oxygen being produced from hydroxide ions:

$4\text{OH}^- \rightarrow \text{O}_2 + 2\text{H}_2\text{O} + 4e^-$4OH−→O2​+2H2​O+4e−

Check each one: the atoms balance, and the total charge is the same on both sides. For $2\text{H}^+ + 2e^- \rightarrow \text{H}_2$2H++2e−→H2​, the left has $2(+1) + 2(-1) = 0$2(+1)+2(−1)=0 and the right is neutral — balanced. Electrons are gained at the cathode (reduction) and lost at the anode (oxidation), so electrolysis is a redox process, linking back to the previous topic.

Exam Tip: A half-equation must balance atoms and charge. Add electrons to the left at the cathode (ions gain electrons) and to the right at the anode (ions lose electrons), and count the charge on each side to check.

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Required Practical: Electrolysis of Aqueous Solutions

A required practical for C3 is to investigate what is produced when aqueous solutions are electrolysed using inert (graphite) electrodes.

Method (numbered):