SC10 · Electrolytic processes · Topic 3 — Chemical changes | GCSE Edexcel Separate Sciences

Revise the key ideas

Electrolysis uses electrical energy from a direct-current (DC) supply to break down (decompose) an electrolyte. It is a chemical change that forms new substances.
Electrolytes are ionic compounds when molten or dissolved in water: their ions are free to move. A solid ionic compound does not conduct because its ions are fixed.
An electrode is a conductor contacting the electrolyte. In an electrolytic cell the cathode is connected to the negative terminal and the anode to the positive terminal.
Positive ions (cations) move to the negative cathode. Negative ions (anions) move to the positive anode.
Electrode signs are for electrolysis driven by a DC supply.
Current in the external metal wires is carried by electrons; in the electrolyte it is carried by moving ions. Do not draw electrons travelling through the solution as the main charge carrier.
Inert electrodes, such as suitable graphite or platinum electrodes, do not react in the GCSE model. Reactive metal electrodes can take part in the reaction and change the products.

OIL RIG means Oxidation Is Loss, Reduction Is Gain of electrons. A half-equation shows the electron transfer for one electrode reaction.
At the cathode, cations can gain electrons and are reduced. Example: Cu²⁺ + 2e⁻ → Cu.
At the anode, anions can lose electrons and are oxidised. Example: 2Br⁻ → Br₂ + 2e⁻.
Balance both atoms and total electric charge. In Cu²⁺ + 2e⁻ → Cu, the left charge +2 − 2 = 0 equals the neutral right side.
Reduction gains electrons at the cathode.
Electrons appear on the left for reduction and on the right for oxidation. A half-equation is not complete if atoms balance but charge does not.
Half-equation writing is Higher-tier material. State symbols may be included when needed: an ion dissolved in water is (aq), not (l).

A molten binary ionic compound contains just two elements and no water. Its own ions form the products: the metal at the cathode and the non-metal at the anode.
Molten lead bromide contains Pb²⁺ and Br⁻. Lead forms at the negative cathode: Pb²⁺ + 2e⁻ → Pb.
Bromine forms at the positive anode: 2Br⁻ → Br₂ + 2e⁻. Overall: PbBr₂ → Pb + Br₂.
Molten PbBr₂ contains Pb²⁺ and Br⁻; water is absent.
Molten sodium chloride gives sodium at the cathode and chlorine at the anode: Na⁺ + e⁻ → Na; 2Cl⁻ → Cl₂ + 2e⁻.
These are predictions or supervised demonstrations, not home experiments: hot melts, reactive metals and harmful halogens require controlled specialist handling.
Do not use the aqueous hydrogen rule for a molten salt; there is no water supplying hydrogen-containing ions.

Water contributes additional possible electrode reactants in aqueous electrolysis, so the products can differ from a molten salt with the same name.
At the cathode, metals less reactive than hydrogen, such as copper, are commonly deposited. If the metal is more reactive than hydrogen, hydrogen is produced instead in the standard GCSE model.
At the anode, concentrated chloride, bromide or iodide solutions commonly produce the corresponding halogen; otherwise oxygen is commonly produced. Conditions, particularly concentration, can affect competition.
Copper chloride solution with sufficient chloride concentration produces copper at the cathode and chlorine at the anode; copper ions gain electrons and chloride ions lose electrons.
Concentrated sodium chloride solution produces hydrogen at the cathode and chlorine at the anode, not sodium metal. Sodium hydroxide remains in the solution.
Sodium sulfate solution with inert electrodes produces hydrogen and oxygen. Water acidified with sulfuric acid also gives hydrogen at the cathode and oxygen at the anode, in a 2:1 gas-volume ratio at the same conditions.
*With sufficient chloride concentration. In sulfate/acidified water, H₂:O₂ volumes are 2:1.
Useful half-equations include 2H⁺ + 2e⁻ → H₂ and 4OH⁻ → O₂ + 2H₂O + 4e⁻, using the ions appropriate to the electrolyte.

Set up copper sulfate solution with inert electrodes connected to a suitable low-voltage DC supply. Keep electrodes separate and use the controlled school procedure with eye protection.
Copper coats the negative cathode: Cu²⁺ + 2e⁻ → Cu. Oxygen bubbles at the positive anode rather than sulfate depositing as a solid.
With inert electrodes, copper ions are removed from solution without a copper anode replacing them. The blue colour becomes paler as Cu²⁺ concentration falls.
With inert electrodes, copper removed from solution is not replenished at the anode.
Identify gases using the teacher's safe method: hydrogen gives a squeaky pop, oxygen relights a glowing splint, and chlorine bleaches damp litmus paper.
Record electrode changes and gas observations. Do not infer product identity from bubbles alone; use the electrode polarity and electrolyte composition.
Electrolysis depends on electrode material as well as solution. Changing from inert to copper electrodes changes what happens at the anode.

With copper electrodes in copper sulfate, copper dissolves from the positive anode: Cu → Cu²⁺ + 2e⁻. The anode loses mass.
Copper ions gain electrons and plate onto the negative cathode: Cu²⁺ + 2e⁻ → Cu. The cathode gains mass.
In the ideal case, anode production of copper ions replaces those removed at the cathode, so copper-ion concentration and blue colour remain approximately constant.
To purify copper, use impure copper as the anode and pure copper as the cathode. Copper transfers through solution onto the cathode; some insoluble impurities collect below the anode as sludge.
Copper ions move through the solution; copper transfers from impure anode to pure cathode.
Compare cleaned, dry electrode masses before and after electrolysis. Control current and time when comparing runs; rinse carefully without removing the deposit and dry consistently.
Plot cathode mass gain against electrolysis time or compare mass changes. Product loss during handling, incomplete drying and side reactions can affect measured results.