ElectrolysisCambridge IGCSE Chemistry: Revision notes
Section 1
What is electrolysis and how does it work?
Electrolysis is the decomposition of an ionic compound when it is molten or in aqueous solution by the passage of an electric current. It is the reverse of a spontaneous redox reaction.
In an electrolytic cell, three key components are required:
- Anode: the positive electrode where oxidation occurs
- Cathode: the negative electrode where reduction occurs
- Electrolyte: the molten or aqueous ionic compound being decomposed
The electrolyte must be molten or aqueous because ions must be able to move freely. Solid ionic compounds cannot undergo electrolysis because their ions are fixed in the crystal lattice and cannot conduct electricity.
During electrolysis:
- Electrons flow from the negative terminal of the power supply through the external circuit to the cathode
- Electrons flow from the anode through the external circuit back to the positive terminal of the power supply
- Inside the electrolyte, anions (negative ions) move towards the anode and cations (positive ions) move towards the cathode
- Ions gain or lose electrons at the electrodes, causing decomposition
Think of electrolysis as 'reversing' a battery: instead of using chemical reactions to make electricity, you're using electricity to break down a compound into its elements.
Section 2
What are the products of electrolysis of molten compounds?
For binary compounds in the molten state (containing only two elements), the products are predictable:
- Metals or hydrogen form at the cathode (reduction: cations gain electrons)
- Non-metals other than hydrogen form at the anode (oxidation: anions lose electrons)
Electrolysis of molten lead(II) bromide (PbBr₂)
| Electrode | Process | Product | Observations |
|---|---|---|---|
| Cathode | Pb²⁺ + 2e⁻ → Pb | Lead (metal) | Shiny grey molten metal forms |
| Anode | 2Br⁻ → Br₂ + 2e⁻ | Bromine (non-metal) | Reddish-brown bromine gas; pungent smell |
Note: Only one product forms at each electrode in molten compounds. The choice of product depends on which ions are present in the molten liquid.
Always write the half-equations separately for each electrode. The examiner wants to see that you understand which species is being oxidised and which is being reduced.
For molten magnesium chloride (MgCl₂): Cathode—Mg²⁺ + 2e⁻ → Mg (shiny metal). Anode—2Cl⁻ → Cl₂ + 2e⁻ (yellow-green toxic gas). Identify the ions present, then apply the rule: metals at cathode, non-metals at anode.
Section 3
What are the products of electrolysis of aqueous solutions?
Electrolysis of aqueous solutions is more complex than molten compounds because water can also be oxidised or reduced. The products depend on the concentration of the electrolyte and the nature of the ions present.
Key rule for aqueous electrolysis:
- At the cathode: Metals less reactive than hydrogen form metallic products; more reactive metals and hydrogen ions cannot compete, so hydrogen gas forms from water reduction: 2H⁺ + 2e⁻ → H₂
- At the anode: Non-metal anions (Br⁻, I⁻, Cl⁻) in concentrated solution are oxidised; in dilute solution or if no suitable anion is present, water is oxidised: 4OH⁻ → O₂ + 2H₂O + 4e⁻ (or 2H₂O → O₂ + 4H⁺ + 4e⁻ in acidic conditions)
Electrolysis of dilute sulfuric acid
| Electrode | Half-equation | Product | Observations |
|---|---|---|---|
| Cathode | 2H⁺ + 2e⁻ → H₂ | Hydrogen gas | Colourless gas; 'pop' with burning splint |
| Anode | 4OH⁻ → O₂ + 2H₂O + 4e⁻ | Oxygen gas | Colourless gas; relights burning splint |
Electrolysis of concentrated aqueous sodium chloride (brine)
| Electrode | Half-equation | Product | Observations |
|---|---|---|---|
| Cathode | 2H⁺ + 2e⁻ → H₂ | Hydrogen gas | Colourless gas |
| Anode | 2Cl⁻ → Cl₂ + 2e⁻ | Chlorine gas | Yellow-green toxic gas; strong smell |
Note: In concentrated NaCl solution, chloride ions are preferentially discharged over hydroxide ions at the anode.
Students often think chloride ions always form chlorine gas at the anode. This is only true in concentrated solution; in dilute solution, water is oxidised to oxygen instead.
Remember: hydrogen gas always forms at the cathode in aqueous solutions (because H⁺ and water reduction are easier than metal ion reduction). At the anode, check the concentration and the nature of the anion to decide between chlorine (or other halogen) and oxygen.
Section 4
How do copper electrodes change the electrolysis of copper(II) sulfate?
When inert electrodes (carbon or platinum) are used with aqueous copper(II) sulfate:
| Electrode | Half-equation | Product |
|---|---|---|
| Cathode | Cu²⁺ + 2e⁻ → Cu | Copper metal (red deposit) |
| Anode | 4OH⁻ → O₂ + 2H₂O + 4e⁻ | Oxygen gas |
When copper electrodes are used instead:
| Electrode | Half-equation | Change |
|---|---|---|
| Cathode | Cu²⁺ + 2e⁻ → Cu | Copper deposits on the electrode (gains mass) |
| Anode | Cu → Cu²⁺ + 2e⁻ | Copper dissolves into solution (loses mass) |
With copper electrodes, the copper anode is sacrificial—it dissolves to replenish the Cu²⁺ ions in solution. The copper cathode grows as Cu²⁺ ions are reduced and deposited. This process is the basis for electrorefining of copper to remove impurities.
Key observations:
- The blue colour of CuSO₄ solution remains constant (Cu²⁺ ions are continuously replenished)
- The cathode gains mass (copper deposited)
- The anode loses mass (copper dissolves)
- Only a small amount of copper is lost as oxygen forms at an inert anode
In the electrolysis of CuSO₄ with inert electrodes: cathode becomes coated with shiny red copper (Cu²⁺ + 2e⁻ → Cu); anode produces colourless oxygen gas (4OH⁻ → O₂ + 2H₂O + 4e⁻). With copper electrodes: cathode gains mass and anode loses mass; the solution colour stays blue because Cu²⁺ is continuously replaced.
Section 5
What is electroplating and how is it used?
Electroplating is the process of coating an object with a thin layer of metal to improve its appearance and resistance to corrosion.
How electroplating is carried out:
- The object to be plated is made the cathode (connected to the negative terminal)
- A piece of the plating metal is made the anode (connected to the positive terminal)
- Both electrodes are immersed in an electrolyte solution containing ions of the plating metal (e.g. a copper salt for copper plating)
- When current flows, metal ions are reduced at the cathode and deposit as a thin coating on the object
- The anode dissolves, replenishing metal ions in the solution
Example: Electroplating with copper
- Object (e.g. steel fork): cathode
- Copper metal: anode
- Electrolyte: copper sulfate solution (CuSO₄)
- Cathode reaction: Cu²⁺ + 2e⁻ → Cu (copper coating forms)
- Anode reaction: Cu → Cu²⁺ + 2e⁻ (copper anode dissolves)
Why electroplating is useful:
- Protects objects from corrosion by preventing oxygen and water from reaching the underlying metal
- Improves appearance (e.g. gold plating on jewellery, chrome plating on car bumpers)
- Makes cheap metals look valuable (e.g. nickel plating steel)
- The coating is thin and uniform, using minimal material
Examiners often ask: 'Describe how to electroplate an object with metal X.' Always state: (1) which electrode is which, (2) what the electrolyte is, (3) what deposits on the object and why, (4) why the anode dissolves.
To silver-plate a brass ring: Connect the ring to the cathode (negative). Connect a silver anode to the positive terminal. Use a silver salt solution (e.g. silver nitrate) as the electrolyte. Ag⁺ + e⁻ → Ag forms a silver coating on the ring; Ag → Ag⁺ + e⁻ dissolves the silver anode.
Section 6
How do ions move and transfer charge during electrolysis?
Charge transfer during electrolysis involves three key processes:
1. Movement of electrons in the external circuit
- Electrons flow from the negative terminal of the power supply through external wires to the cathode
- Electrons are released at the anode and flow back through external wires to the positive terminal of the power supply
- This flow of electrons constitutes the electric current
2. Gain and loss of electrons at the electrodes
- At the cathode (reduction): Cations are reduced and gain electrons. For example: Cu²⁺ + 2e⁻ → Cu
- At the anode (oxidation): Anions are oxidised and lose electrons. For example: 2Cl⁻ → Cl₂ + 2e⁻
- The number of electrons gained at the cathode equals the number lost at the anode (electrons are conserved)
3. Movement of ions in the electrolyte
- Cations (positive ions) move towards the cathode (negative electrode) and are reduced there
- Anions (negative ions) move towards the anode (positive electrode) and are oxidised there
- This movement of charged particles (ions) within the electrolyte allows charge to be transferred and completes the circuit
Constructing ionic half-equations:
Half-equations show the transfer of electrons at each electrode. They must be balanced for both atoms and charge.
Reduction half-equation (cathode):
- Cation gains electrons
- Example: Cu²⁺ + 2e⁻ → Cu (2 electrons gained; atom balanced; charge balanced: 2+ − 2 = 0)
Oxidation half-equation (anode):
- Anion loses electrons
- Example: 2Cl⁻ → Cl₂ + 2e⁻ (2 electrons lost; atoms balanced; charge balanced: 2 − − 2 = 0)
Writing half-equations for halide compounds in aqueous solution:
| Halide | Concentration | Anode product | Half-equation |
|---|---|---|---|
| Cl⁻ | Concentrated | Cl₂ | 2Cl⁻ → Cl₂ + 2e⁻ |
| Cl⁻ | Dilute | O₂ | 4OH⁻ → O₂ + 2H₂O + 4e⁻ |
| Br⁻ | Concentrated | Br₂ | 2Br⁻ → Br₂ + 2e⁻ |
| Br⁻ | Dilute | O₂ | 4OH⁻ → O₂ + 2H₂O + 4e⁻ |
| I⁻ | Any | I₂ | 2I⁻ → I₂ + 2e⁻ |
Note: Iodide ions are always oxidised to iodine (never to oxygen) because iodide is the most easily oxidised halide.
When writing half-equations, always check: (1) Are atoms balanced? (2) Are charges balanced? (3) Are electrons shown on the correct side? Reduction has e⁻ on the left; oxidation has e⁻ on the right.
For electrolysis of concentrated KBr solution: Cathode: 2H⁺ + 2e⁻ → H₂. Anode: 2Br⁻ → Br₂ + 2e⁻. Check—atoms balanced? Yes (2 H, 2 Br). Charges balanced? Cathode: 2(+1) + 2(−1) = 0 on left; 0 on right. Anode: 2(−1) = −2; Br₂ + 2(−1) = −2. Electrons equal on both sides.
Students often forget that at the cathode, it is usually water being reduced to hydrogen (not a metal ion), especially in dilute aqueous solutions. Write 2H⁺ + 2e⁻ → H₂ or 2H₂O + 2e⁻ → H₂ + 2OH⁻, not a metal ion reduction.
Must Know
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Electrolysis is the decomposition of an ionic compound (molten or aqueous) by electric current. In an electrolytic cell, the anode is positive, the cathode is negative, and the electrolyte conducts ions.
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At the cathode: metals or hydrogen form (reduction, cations gain electrons). At the anode: non-metals other than hydrogen form (oxidation, anions lose electrons).
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For molten binary compounds, products are predictable: metals at cathode, non-metals at anode. For example, molten PbBr₂ gives Pb (cathode) and Br₂ (anode).
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For aqueous electrolysis, the product at the anode depends on concentration: concentrated halide solutions give halogen (Cl₂, Br₂, I₂); dilute solutions or no halide give oxygen. At the cathode, hydrogen always forms (water is reduced).
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Copper electrodes dissolve during electrolysis: the anode dissolves to replenish Cu²⁺ ions; the cathode gains mass as copper deposits. With inert electrodes, oxygen forms at the anode instead.
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Electroplating: the object is the cathode, the plating metal is the anode (dissolves), and metal ions from the electrolyte deposit on the object, improving appearance and corrosion resistance.
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Charge transfer: electrons flow through the external circuit (cathode to anode through power supply); cations move to the cathode and gain electrons (reduction); anions move to the anode and lose electrons (oxidation). Half-equations must show electron transfer and be balanced for atoms and charge.
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Halide predictions: In concentrated solution, halides (Cl⁻, Br⁻, I⁻) are oxidised to halogens. In dilute solution, water is oxidised to O₂ instead. Iodide is always oxidised to I₂ (never to oxygen).
That's the notes covered.
Carry on to the next subtopic.