Chemical Changes

What you'll learn

Chemical Changes covers fundamental reactions including the reactivity series of metals, acids, alkalis and neutralisation, as well as electrolysis and redox processes. This topic forms approximately 18% of your Edexcel GCSE Chemistry examination and includes both paper 1 and paper 2 content. Understanding these concepts is essential for both Foundation and Higher tier students.

Key terms and definitions

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

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

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

Spectator ions — ions that are present in a reaction mixture but do not participate in the chemical reaction

Neutralisation — a reaction between an acid and a base to produce a salt and water, with pH moving towards 7

Strong acid — an acid that completely dissociates (ionises) into ions in aqueous solution

Core concepts

The reactivity series and metal extraction

The reactivity series ranks metals in order of their reactivity. From most reactive to least reactive:

Potassium, Sodium, Lithium, Calcium, Magnesium, Aluminium, (Carbon), Zinc, Iron, (Hydrogen), Copper, Silver, Gold

Carbon and hydrogen are non-metals included for comparison when extracting metals from their ores.

Metal extraction methods:

• Metals more reactive than carbon (potassium to aluminium) must be extracted by electrolysis of molten compounds — this is expensive due to high energy costs
• Metals less reactive than carbon (zinc to iron) can be extracted by reduction with carbon in a displacement reaction
• Metals less reactive than hydrogen (copper to gold) may be found native (uncombined) or extracted by heating alone

Displacement reactions occur when a more reactive metal displaces a less reactive metal from its compound:

magnesium + copper sulfate → magnesium sulfate + copper

Mg(s) + CuSO₄(aq) → MgSO₄(aq) + Cu(s)

This works because magnesium is higher in the reactivity series than copper.

Reactions of metals with acids and water

Metals with dilute acids:

Only metals above hydrogen in the reactivity series react with dilute acids to produce a salt and hydrogen gas:

metal + acid → salt + hydrogen

For example: zinc + hydrochloric acid → zinc chloride + hydrogen

Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)

Metals with water:

• Very reactive metals (potassium, sodium, lithium, calcium) react vigorously with cold water to produce metal hydroxide and hydrogen
• Less reactive metals (magnesium, zinc, iron) react slowly with water but react with steam to produce metal oxide and hydrogen

magnesium + steam → magnesium oxide + hydrogen

Mg(s) + H₂O(g) → MgO(s) + H₂(g)

Acids, bases and neutralisation

Acids produce hydrogen ions (H⁺) in aqueous solution. The pH scale measures acidity from 0-14:

• pH 0-6: acidic
• pH 7: neutral
• pH 8-14: alkaline

Strong acids completely ionise in water:

• Hydrochloric acid (HCl)
• Sulfuric acid (H₂SO₄)
• Nitric acid (HNO₃)

Weak acids only partially ionise in water:

• Ethanoic acid (CH₃COOH)
• Citric acid
• Carbonic acid

For equal concentrations, strong acids have lower pH and react faster than weak acids.

Bases are substances that neutralise acids. Alkalis are soluble bases that produce hydroxide ions (OH⁻) in water.

Neutralisation reactions:

acid + base → salt + water

The general ionic equation for neutralisation is:

H⁺(aq) + OH⁻(aq) → H₂O(l)

Specific reactions of acids:

1. Acid + metal oxide/hydroxide → salt + water

   sulfuric acid + copper oxide → copper sulfate + water

   H₂SO₄(aq) + CuO(s) → CuSO₄(aq) + H₂O(l)

2. Acid + metal carbonate → salt + water + carbon dioxide

   hydrochloric acid + calcium carbonate → calcium chloride + water + carbon dioxide

   2HCl(aq) + CaCO₃(s) → CaCl₂(aq) + H₂O(l) + CO₂(g)

Naming salts:

The salt name comes from:

• First part: the metal (or ammonium)
• Second part: the acid used
  • Hydrochloric acid → chloride
  • Sulfuric acid → sulfate
  • Nitric acid → nitrate

Electrolysis fundamentals

Electrolysis uses electrical energy to decompose ionic compounds. Key components:

• Electrodes — conductors through which current enters and leaves the electrolyte
• Anode — positive electrode where oxidation occurs (anions attracted here)
• Cathode — negative electrode where reduction occurs (cations attracted here)

At the cathode (negative electrode):

• Positive ions (cations) are attracted
• Cations gain electrons (reduction)
• Metals or hydrogen are produced

At the anode (positive electrode):

• Negative ions (anions) are attracted
• Anions lose electrons (oxidation)
• Non-metals (often oxygen or halogens) are produced

Half equations show electron transfer at each electrode:

At cathode: Cu²⁺ + 2e⁻ → Cu (reduction)

At anode: 2O²⁻ → O₂ + 4e⁻ (oxidation)

Electrolysis of molten ionic compounds

When molten, ionic compounds contain free-moving ions. Electrolysis produces elements:

Lead bromide (molten):

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

The compound always splits into its constituent elements when molten.

Electrolysis of aqueous solutions

Aqueous solutions contain water molecules that can also be discharged. Rules for predicting products:

At the cathode:

• If the metal is less reactive than hydrogen (below hydrogen in reactivity series), the metal is produced
• If the metal is more reactive than hydrogen, hydrogen gas is produced instead

At the anode:

• If a halide ion (Cl⁻, Br⁻, I⁻) is present, the halogen is produced
• Otherwise, oxygen gas is produced from hydroxide ions

Example: Copper sulfate solution:

• Ions present: Cu²⁺, H⁺, SO₄²⁻, OH⁻
• Cathode: Cu²⁺ + 2e⁻ → Cu (copper less reactive than hydrogen)
• Anode: 4OH⁻ → O₂ + 2H₂O + 4e⁻ (no halide present)

Example: Sodium chloride solution (brine):

• Ions present: Na⁺, H⁺, Cl⁻, OH⁻
• Cathode: 2H⁺ + 2e⁻ → H₂ (sodium more reactive than hydrogen)
• Anode: 2Cl⁻ → Cl₂ + 2e⁻ (halide present)

This is industrially important for producing chlorine, hydrogen and sodium hydroxide solution.

Oxidation and reduction (redox)

Redox reactions involve the transfer of electrons. Remember OILRIG:

• Oxidation Is Loss (of electrons)
• Reduction Is Gain (of electrons)

Identifying oxidation and reduction:

In the reaction: Mg + CuO → MgO + Cu

• Magnesium loses electrons: Mg → Mg²⁺ + 2e⁻ (oxidation)
• Copper gains electrons: Cu²⁺ + 2e⁻ → Cu (reduction)

Copper oxide is the oxidising agent (it oxidises magnesium). Magnesium is the reducing agent (it reduces copper).

Alternative definitions:

Oxidation can also be defined as:

• Gain of oxygen
• Loss of hydrogen

Reduction can also be defined as:

• Loss of oxygen
• Gain of hydrogen

These definitions apply when electron transfer isn't obvious.

Worked examples

Example 1: Writing ionic equations

Question: Write a balanced symbol equation and ionic equation for the reaction between zinc and copper sulfate solution. Identify the spectator ions. [4 marks]

Answer:

Symbol equation: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s) [1 mark]

Full ionic equation: Zn(s) + Cu²⁺(aq) + SO₄²⁻(aq) → Zn²⁺(aq) + SO₄²⁻(aq) + Cu(s) [1 mark]

Ionic equation (removing spectator ions): Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s) [1 mark]

Spectator ions: SO₄²⁻ [1 mark]

Example 2: Predicting electrolysis products

Question: A student electrolyses dilute sulfuric acid using inert electrodes.

(a) Name the products at the cathode and anode. [2 marks]

(b) Write half equations for the reactions at each electrode. [2 marks]

Answer:

(a) Cathode: hydrogen gas [1 mark] Anode: oxygen gas [1 mark]

(b) Cathode: 2H⁺ + 2e⁻ → H₂ [1 mark] Anode: 4OH⁻ → O₂ + 2H₂O + 4e⁻ [1 mark]

Example 3: Neutralisation and salt preparation

Question: A student makes copper sulfate crystals by adding excess copper oxide to warm dilute sulfuric acid.

(a) Write a balanced symbol equation for this reaction. [2 marks]

(b) Explain why excess copper oxide is used. [1 mark]

(c) Describe how the student would obtain pure, dry copper sulfate crystals from the reaction mixture. [3 marks]

Answer:

(a) CuO(s) + H₂SO₄(aq) → CuSO₄(aq) + H₂O(l) [2 marks — correct formulae and balanced]

(b) To ensure all the acid has reacted / to make sure no acid remains [1 mark]

(c)

• Filter to remove excess copper oxide [1 mark]
• Heat the filtrate to evaporate some water / concentrate the solution [1 mark]
• Leave to crystallise / cool slowly to form crystals [1 mark]
• Pat dry with filter paper / dry in warm oven [1 mark for any third point]

Common mistakes and how to avoid them

• Confusing oxidation and reduction — Use OILRIG consistently. Remember that in electrolysis, oxidation always occurs at the anode, reduction at the cathode (think "vowels together": Anode-Oxidation)

• Incorrect salt names — The salt name must match the acid used: hydrochloric → chloride, sulfuric → sulfate, nitric → nitrate. Students often write "sulfide" instead of "sulfate"

• Predicting wrong products in aqueous electrolysis — Apply the rules systematically: check the metal's reactivity for cathode products; check for halides for anode products. Don't assume the compound simply splits

• Unbalanced half equations — Always ensure charge is balanced as well as atoms. Add electrons to balance the charge, then check total charge is equal on both sides

• Writing spectator ions in ionic equations — Ionic equations only show species that change. Remove all ions that appear unchanged on both sides of the equation

• pH scale confusion — Remember acids are pH 0-6 (lower numbers = more acidic), neutral is pH 7, and alkalis are pH 8-14 (higher numbers = more alkaline). Strong acids are not "pH 1" — concentration matters too

Exam technique for "Chemical Changes"

• Command words matter — "State" requires a simple answer with no explanation. "Explain" requires a reason using scientific terminology. "Describe" needs a clear account without necessarily explaining why

• Half equations in electrolysis — Always show state symbols, balance atoms first, then add electrons to balance charge. For Higher tier, you must be able to write these from memory for common ions

• Show your working — In calculation questions (e.g., involving moles and electrolysis), write each step clearly. Marks are often awarded for method even if the final answer is incorrect

• 6-mark questions — Structure extended answers logically with clear scientific terminology. Cover multiple points with detail. Link ideas together using words like "therefore," "because," and "this means that"

Quick revision summary

Chemical Changes covers reactivity of metals (extraction methods and displacement reactions), acid-base chemistry (neutralisation, pH, salt preparation), and electrolysis (molten and aqueous). Master the reactivity series for predicting reactions. Understand that strong acids completely ionise while weak acids partially ionise. In electrolysis, reduction occurs at the cathode (gain electrons) and oxidation at the anode (lose electrons). For aqueous electrolysis, apply rules based on reactivity and ion type. Remember OILRIG for redox reactions and practice writing balanced half equations.