Chemical Changes

What you'll learn

Chemical Changes forms a substantial component of Edexcel GCSE Chemistry, testing your understanding of reactivity, extraction processes, acid-base reactions, and electrolysis. This topic appears across Paper 1 and synoptic questions in Paper 2, typically accounting for 15-20% of available marks. Mastery of this content requires both theoretical knowledge and the ability to apply concepts to unfamiliar contexts.

Key terms and definitions

Oxidation — the loss of electrons, gain of oxygen, or loss of hydrogen by a substance during a chemical reaction.

Reduction — the gain of electrons, loss of oxygen, or gain of hydrogen by a substance during a chemical reaction.

Redox reaction — a chemical reaction in which both oxidation and reduction occur simultaneously.

Displacement reaction — a reaction where a more reactive element takes the place of a less reactive element in a compound.

Electrolysis — the decomposition of an ionic compound when molten or in aqueous solution by passing an electric current through it.

Electrolyte — a substance that conducts electricity when molten or dissolved in water and is decomposed during electrolysis.

Neutralisation — a reaction between an acid and a base that produces a salt and water only.

Salt — an ionic compound formed when the hydrogen ion in an acid is replaced by a metal ion or ammonium ion.

Core concepts

The reactivity series and displacement reactions

The reactivity series ranks metals in order of their reactivity, from most reactive to least reactive. For Edexcel GCSE Chemistry, you must know this sequence:

Potassium > Sodium > Lithium > Calcium > Magnesium > Aluminium > Carbon > Zinc > Iron > Hydrogen > Copper > Silver > Gold

Note that carbon and hydrogen are non-metals included for comparison purposes, particularly in extraction chemistry.

A more reactive metal will displace a less reactive metal from its compound. This principle underpins displacement reactions:

• Magnesium + copper sulfate → magnesium sulfate + copper
• Zinc + lead nitrate → zinc nitrate + lead
• Iron + copper chloride → iron chloride + copper

No reaction occurs when a less reactive metal is added to a compound of a more reactive metal. For example, copper cannot displace zinc from zinc sulfate.

When metals react with water or dilute acids, reactivity determines the vigour of the reaction:

• Very reactive metals (potassium to calcium) react vigorously with cold water
• Moderately reactive metals (magnesium to iron) react with steam or acids
• Unreactive metals (below hydrogen) do not react with water or dilute acids

Extraction of metals and reduction

The position of a metal in the reactivity series determines the extraction method required.

Metals more reactive than carbon (potassium to aluminium) must be extracted by electrolysis of molten compounds. This is expensive due to high energy requirements. Aluminium extraction uses the Hall-Héroult process:

• Aluminium oxide (from bauxite ore) is dissolved in molten cryolite
• Temperature reduced from 2050°C to approximately 950°C
• At the cathode: Al³⁺ + 3e⁻ → Al (reduction)
• At the anode: 2O²⁻ → O₂ + 4e⁻ (oxidation)
• Carbon anodes must be replaced regularly as they react with oxygen forming carbon dioxide

Metals less reactive than carbon (zinc to silver) can be extracted by reduction with carbon in a blast furnace or similar process. Iron extraction in the blast furnace:

1. Iron ore (haematite, Fe₂O₃), coke (carbon), and limestone (CaCO₃) added at the top
2. Hot air blasted in at the bottom
3. Carbon burns: C + O₂ → CO₂
4. Carbon dioxide reduced: CO₂ + C → 2CO
5. Iron oxide reduced: Fe₂O₃ + 3CO → 2Fe + 3CO₂
6. Molten iron tapped off at the bottom
7. Limestone removes impurities as slag: CaCO₃ → CaO + CO₂, then CaO + SiO₂ → CaSiO₃

Unreactive metals (below hydrogen) occur naturally as elements and require no extraction.

Oxidation and reduction in terms of electrons

In Edexcel GCSE Chemistry, oxidation and reduction must be understood at the electron transfer level, particularly for electrolysis questions.

OIL RIG is a useful mnemonic: Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons).

In the reaction: Mg + CuSO₄ → MgSO₄ + Cu

• Magnesium atoms lose electrons: Mg → Mg²⁺ + 2e⁻ (oxidation)
• Copper ions gain electrons: Cu²⁺ + 2e⁻ → Cu (reduction)
• Magnesium is oxidised; copper ions are reduced
• This is a redox reaction

Identifying oxidation and reduction in ionic equations is essential for higher-tier papers.

Reactions of acids

Acids produce hydrogen ions (H⁺) in aqueous solutions. The three main acids at GCSE are:

• Hydrochloric acid (HCl) → produces chloride salts
• Sulfuric acid (H₂SO₄) → produces sulfate salts
• Nitric acid (HNO₃) → produces nitrate salts

Bases are substances that neutralise acids. Metal oxides and metal hydroxides are bases. Bases that dissolve in water are called alkalis and produce hydroxide ions (OH⁻).

The pH scale measures acidity from 0-14, with 7 being neutral. Acids have pH <7, alkalis pH >7.

Key acid reactions tested on Edexcel papers:

Acid + metal → salt + hydrogen

• Magnesium + sulfuric acid → magnesium sulfate + hydrogen
• 2HCl + Mg → MgCl₂ + H₂

Acid + metal oxide → salt + water

• Copper oxide + sulfuric acid → copper sulfate + water
• CuO + H₂SO₄ → CuSO₄ + H₂O

Acid + metal hydroxide → salt + water (neutralisation)

• Sodium hydroxide + hydrochloric acid → sodium chloride + water
• NaOH + HCl → NaCl + H₂O

Acid + metal carbonate → salt + water + carbon dioxide

• Calcium carbonate + nitric acid → calcium nitrate + water + carbon dioxide
• CaCO₃ + 2HNO₃ → Ca(NO₃)₂ + H₂O + CO₂

The general ionic equation for neutralisation is: H⁺(aq) + OH⁻(aq) → H₂O(l)

Making salts

Soluble salts can be prepared by reacting acids with excess solid (metal, oxide, hydroxide, or carbonate):

1. Add excess solid to warm dilute acid until no more reacts
2. Filter to remove unreacted solid
3. Evaporate solution to crystallisation point
4. Leave to crystallise
5. Dry crystals with filter paper

Insoluble salts are prepared by precipitation:

1. Mix solutions containing the required ions
2. Filter to collect the precipitate
3. Wash with distilled water
4. Dry the precipitate

Example: Making lead iodide (insoluble)

• Lead nitrate + potassium iodide → lead iodide + potassium nitrate
• Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq)

Electrolysis principles and applications

Electrolysis splits ionic compounds using electricity. The compound must be molten or dissolved so ions are free to move.

Key components:

• Cathode — negative electrode where reduction occurs (positive ions gain electrons)
• Anode — positive electrode where oxidation occurs (negative ions lose electrons)

Electrolysis of molten ionic compounds:

For molten lead bromide (PbBr₂):

• At cathode: Pb²⁺ + 2e⁻ → Pb (lead metal forms)
• At anode: 2Br⁻ → Br₂ + 2e⁻ (bromine gas forms)

Electrolysis of aqueous solutions:

More complex because water molecules (H₂O) can also be discharged, producing H₂ at the cathode or O₂ at the anode.

Rules for aqueous electrolysis:

At the cathode (negative electrode):

• If the metal is less reactive than hydrogen, the metal is produced
• If the metal is more reactive than hydrogen, hydrogen gas is produced

At the anode (positive electrode):

• If halide ions (Cl⁻, Br⁻, I⁻) are present, the halogen is produced
• If no halide ions, or with dilute solutions, oxygen is produced from OH⁻ ions

Electrolysis of copper sulfate solution with inert electrodes:

• Cathode: Cu²⁺ + 2e⁻ → Cu (copper deposited)
• Anode: 4OH⁻ → O₂ + 2H₂O + 4e⁻ (oxygen gas produced)

Electrolysis of sodium chloride solution (brine):

• Cathode: 2H⁺ + 2e⁻ → H₂ (hydrogen gas, as sodium is more reactive)
• Anode: 2Cl⁻ → Cl₂ + 2e⁻ (chlorine gas, halide present)
• Sodium hydroxide solution remains (industrial chlor-alkali process)

Electroplating uses electrolysis to coat objects with a thin metal layer. The object is the cathode, and metal ions in solution are deposited onto it.

Worked examples

Example 1: Displacement reaction

Question: A student adds iron filings to a solution of copper sulfate. A reaction occurs, forming copper metal and a green solution.

(a) Write a word equation for this reaction. [1 mark]

(b) Explain why this reaction occurs in terms of reactivity. [2 marks]

(c) Write a balanced symbol equation including state symbols. [3 marks]

Answers:

(a) Iron + copper sulfate → iron sulfate + copper

(b) Iron is more reactive than copper [1 mark]. Therefore, iron displaces copper from copper sulfate solution [1 mark].

(c) Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s) [2 marks for correct equation, 1 mark for state symbols]

Example 2: Electrolysis half-equations

Question: Aluminium is extracted from aluminium oxide by electrolysis.

(a) Write a half-equation for the reaction at the negative electrode. [2 marks]

(b) Explain why the carbon anodes must be regularly replaced. [2 marks]

Answers:

(a) Al³⁺ + 3e⁻ → Al [1 mark for correct formula, 1 mark for balanced electrons]

(b) Oxygen is produced at the anode [1 mark]. This reacts with the carbon anode to form carbon dioxide, wearing the anode away [1 mark].

Example 3: Preparing a salt

Question: A student prepares copper sulfate crystals by reacting excess copper oxide with dilute sulfuric acid.

(a) How can the student tell when all the acid has reacted? [1 mark]

(b) Describe how the student obtains pure, dry crystals from the mixture. [3 marks]

Answers:

(a) Copper oxide will remain at the bottom/excess solid visible/no more dissolves [1 mark]

(b) Filter to remove excess copper oxide [1 mark]. Heat the solution to evaporate water until crystals start to form/crystallisation point reached [1 mark]. Leave to crystallise, then dry with filter paper [1 mark].

Common mistakes and how to avoid them

• Confusing the reactivity series order. Students often misplace aluminium or reverse iron and zinc. Learn the mnemonic "Please Send Lions, Cats, Monkeys And Cute Zebras Into Hot Countries Signed Gordon" or create your own. Carbon sits between aluminium and zinc.

• Writing incorrect salt names. The salt name comes from the metal (or ammonium) plus the acid: hydrochloric acid makes chlorides, sulfuric acid makes sulfates, nitric acid makes nitrates. Zinc oxide + nitric acid makes zinc nitrate, not "zinc nitric" or "zinc oxide nitrate."

• Mixing up cathode and anode. Cathode is negative, anode is positive. Use "CAthode = Negative = CATION attracted" to remember that positive cations move to the negative cathode. Reduction happens at the cathode (both start with consonants).

• Forgetting state symbols in electrolysis equations. Edexcel mark schemes often allocate a mark specifically for state symbols: (s), (l), (aq), (g). Electrodes produce elements as (s) for metals or (g) for gases. Solutions are (aq).

• Stating that unreactive metals are extracted by reduction. Metals below hydrogen in the reactivity series (copper, silver, gold) are found naturally as elements and need no extraction process. They are unreactive enough to exist uncombined.

• Not balancing electrons in half-equations. Each half-equation must show electrons explicitly. The number of electrons lost in oxidation must equal the number gained in reduction. For Al³⁺ + e⁻ → Al, you must write 3e⁻, not e⁻.

Exam technique for Chemical Changes

• "Explain" questions require because/therefore statements. For 2-mark "explain" questions on reactivity or displacement, state the relevant position in the reactivity series, then link this to what happens using "therefore" or "so". Example: "Magnesium is more reactive than zinc, therefore magnesium displaces zinc from zinc sulfate."

• Half-equations must show charges and electrons explicitly. Both sides of the equation must balance for atoms and charges. Write the electrons on the correct side: gain electrons (reduction) means electrons on the left of the arrow; loss of electrons (oxidation) means electrons on the right.

• Method questions follow logical sequence. When describing salt preparation or electrolysis setup, use numbered points or sequence words (first, then, next, finally). Include safety precautions where relevant for full marks.

• Learn the general equations for acid reactions. Pattern recognition saves time. All acid + carbonate reactions produce salt + water + carbon dioxide. Apply the formula rather than memorising individual reactions.

Quick revision summary

The reactivity series determines displacement reactions and extraction methods. More reactive metals displace less reactive ones from compounds. Metals above carbon need electrolysis; those below carbon are reduced with carbon or occur naturally. Oxidation is electron loss; reduction is electron gain. Acids react with metals, bases, and carbonates to form salts. Neutralisation produces salt and water only. Electrolysis decomposes molten or dissolved ionic compounds. At the cathode, reduction occurs (positive ions gain electrons); at the anode, oxidation occurs (negative ions lose electrons). Half-equations must balance atoms and charges with electrons shown explicitly.