What you'll learn
This topic focuses on your ability to predict what happens when chemicals react and identify the products formed. You'll learn to recognise patterns in chemical reactions, use reactivity series to predict outcomes, and apply chemical tests to identify unknown substances. These skills are fundamental to both exam success and practical chemistry.
Key terms and definitions
Reactivity series — an ordered list of metals (and carbon/hydrogen) arranged by their tendency to lose electrons and form positive ions, with most reactive at the top
Displacement reaction — a reaction where a more reactive element takes the place of a less reactive element in a compound
Spectator ions — ions present in a reaction mixture that do not participate in the actual chemical change and remain unchanged
Precipitate — an insoluble solid formed when two solutions react together
Oxidation — the loss of electrons, or gain of oxygen, or loss of hydrogen
Reduction — the gain of electrons, or loss of oxygen, or gain of hydrogen
Neutralisation — a reaction between an acid and a base that produces a salt and water, with pH moving towards 7
Functional group — a specific group of atoms within a molecule that determines the characteristic chemical reactions of that molecule
Core concepts
Types of chemical reactions you must recognise
At GCSE level, you need to identify and predict several key reaction types:
Combustion reactions occur when substances react with oxygen, releasing energy. Complete combustion of hydrocarbons produces carbon dioxide and water:
- methane + oxygen → carbon dioxide + water
- CH₄ + 2O₂ → CO₂ + 2H₂O
Incomplete combustion occurs with limited oxygen, producing carbon monoxide or carbon (soot).
Thermal decomposition involves breaking down compounds using heat:
- metal carbonates → metal oxide + carbon dioxide
- CaCO₃ → CaO + CO₂
Oxidation and reduction (redox) reactions involve electron transfer. Remember "OIL RIG": Oxidation Is Loss, Reduction Is Gain (of electrons).
In the reaction: magnesium + copper sulfate → magnesium sulfate + copper
- Magnesium loses electrons (oxidised): Mg → Mg²⁺ + 2e⁻
- Copper ions gain electrons (reduced): Cu²⁺ + 2e⁻ → Cu
Using the reactivity series to predict reactions
The reactivity series (from most to least reactive) is: Potassium, Sodium, Lithium, Calcium, Magnesium, Aluminium, (Carbon), Zinc, Iron, (Hydrogen), Copper, Silver, Gold
Memorise using: "Please Send Lions, Cats, Monkeys And Cute Zebras Into Hot Countries Signed Gold"
Key predictions using the series:
Metal + acid reactions: Only metals above hydrogen react with acids to produce salt + hydrogen gas
- Zinc reacts with hydrochloric acid: Zn + 2HCl → ZnCl₂ + H₂
- Copper does NOT react with dilute acids (below hydrogen)
Metal + water reactions: Only very reactive metals (potassium to calcium) react vigorously with cold water
- Sodium + water → sodium hydroxide + hydrogen
- 2Na + 2H₂O → 2NaOH + H₂
- Less reactive metals like iron only react with steam
Displacement reactions: A more reactive metal displaces a less reactive metal from its compound
- Iron + copper sulfate → iron sulfate + copper (iron more reactive than copper)
- Fe + CuSO₄ → FeSO₄ + Cu
- Copper + magnesium sulfate → NO REACTION (copper less reactive than magnesium)
Extraction of metals: Metals below carbon in the series can be extracted by reduction with carbon; metals above carbon require electrolysis
Predicting products from acid-base reactions
Acids react with bases to produce salts through neutralisation. The salt name depends on the acid and base used:
Hydrochloric acid (HCl) produces chloride salts Sulfuric acid (H₂SO₄) produces sulfate salts Nitric acid (HNO₃) produces nitrate salts
The metal part of the salt comes from the base:
Metal oxides + acids:
- copper oxide + sulfuric acid → copper sulfate + water
- CuO + H₂SO₄ → CuSO₄ + H₂O
Metal hydroxides + acids:
- sodium hydroxide + hydrochloric acid → sodium chloride + water
- NaOH + HCl → NaCl + H₂O
Metal carbonates + acids:
- calcium carbonate + nitric acid → calcium nitrate + water + carbon dioxide
- CaCO₃ + 2HNO₃ → Ca(NO₃)₂ + H₂O + CO₂
Ammonia + acids (ammonia acts as a base):
- ammonia + sulfuric acid → ammonium sulfate
- 2NH₃ + H₂SO₄ → (NH₄)₂SO₄
Precipitation reactions and identifying ions
Precipitation occurs when two soluble salts react to form an insoluble product. You can predict precipitates using solubility rules:
Most chlorides, bromides, and iodides are soluble EXCEPT silver and lead halides Most sulfates are soluble EXCEPT barium sulfate, calcium sulfate, and lead sulfate Most carbonates are insoluble EXCEPT Group 1 carbonates and ammonium carbonate Most hydroxides are insoluble EXCEPT Group 1 hydroxides and barium hydroxide
Testing for cations (positive ions):
Flame tests identify metal ions by characteristic colours:
- Lithium → crimson/red
- Sodium → yellow/orange
- Potassium → lilac/purple
- Calcium → orange-red
- Copper → blue-green
Sodium hydroxide solution added to solutions containing metal ions:
- Calcium: white precipitate
- Copper(II): blue precipitate
- Iron(II): green precipitate
- Iron(III): brown precipitate
- Aluminium: white precipitate (dissolves in excess NaOH)
- Magnesium: white precipitate
Testing for anions (negative ions):
Carbonate test: Add dilute acid → effervescence (fizzing) produces CO₂, which turns limewater milky
Halide test: Add dilute nitric acid then silver nitrate solution:
- Chloride ions → white precipitate (AgCl)
- Bromide ions → cream precipitate (AgBr)
- Iodide ions → yellow precipitate (AgI)
Sulfate test: Add dilute hydrochloric acid then barium chloride solution:
- Sulfate ions → white precipitate (BaSO₄)
Writing ionic equations
Full equations show all substances, but ionic equations show only the particles that actually change in the reaction.
For the reaction: silver nitrate + sodium chloride → silver chloride + sodium nitrate
Full equation: AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)
Full ionic equation: Ag⁺(aq) + NO₃⁻(aq) + Na⁺(aq) + Cl⁻(aq) → AgCl(s) + Na⁺(aq) + NO₃⁻(aq)
Net ionic equation (removing spectator ions Na⁺ and NO₃⁻): Ag⁺(aq) + Cl⁻(aq) → AgCl(s)
This shows that any soluble silver salt reacting with any soluble chloride will produce the same precipitate.
Predicting organic reaction products
Combustion of alcohols: Alcohols burn in oxygen to produce carbon dioxide and water:
- ethanol + oxygen → carbon dioxide + water
- C₂H₅OH + 3O₂ → 2CO₂ + 3H₂O
Oxidation of alcohols: Alcohols can be oxidised to carboxylic acids using oxidising agents:
- ethanol → ethanoic acid
- C₂H₅OH + [O] → CH₃COOH + H₂O (where [O] represents an oxidising agent)
Reaction of carboxylic acids: Carboxylic acids react like other acids:
- With carbonates: ethanoic acid + sodium carbonate → sodium ethanoate + water + carbon dioxide
- 2CH₃COOH + Na₂CO₃ → 2CH₃COONa + H₂O + CO₂
Worked examples
Example 1: Predicting displacement reactions (3 marks)
Question: A student adds iron powder to a solution of copper sulfate. Predict what would happen and write a balanced symbol equation for any reaction that occurs. Explain your answer using the reactivity series.
Answer: A reaction would occur because iron is more reactive than copper. [1 mark]
The blue copper sulfate solution would fade and brown copper metal would form. [1 mark]
Fe + CuSO₄ → FeSO₄ + Cu [1 mark]
Examiner note: You must reference the reactivity series to explain WHY the reaction occurs.
Example 2: Identifying unknown ions (4 marks)
Question: A student has a solution of an unknown salt. They carry out the following tests:
Test 1: Flame test produces a lilac colour Test 2: Adding dilute nitric acid then silver nitrate produces a cream precipitate
Identify the salt and explain your reasoning.
Answer: The lilac flame test indicates potassium ions (K⁺) are present. [1 mark]
The cream precipitate with silver nitrate indicates bromide ions (Br⁻) are present. [1 mark]
Therefore, the salt is potassium bromide. [1 mark]
The precipitate is silver bromide: Ag⁺(aq) + Br⁻(aq) → AgBr(s) [1 mark]
Examiner note: Always state both the test result AND what it indicates.
Example 3: Predicting salt formation (2 marks)
Question: Name the salt formed when calcium hydroxide reacts with nitric acid. Write the word equation for this reaction.
Answer: The salt is calcium nitrate. [1 mark]
calcium hydroxide + nitric acid → calcium nitrate + water [1 mark]
Examiner note: Remember the metal/base provides the first part of the salt name, the acid provides the second part.
Common mistakes and how to avoid them
Confusing reactivity with reaction speed: A more reactive metal doesn't always react faster visibly. Aluminium is very reactive but has a protective oxide layer that slows initial reaction with water.
Incorrect salt names: Students often write "calcium chlorate" instead of "calcium chloride" when hydrochloric acid is used. Remember: hydrochloric → chloride, sulfuric → sulfate, nitric → nitrate.
Forgetting state symbols in ionic equations: Precipitates must be shown as (s), ions in solution as (aq). Writing AgCl(aq) instead of AgCl(s) loses marks.
Not removing spectator ions properly: When writing ionic equations, you must show the same spectator ions on both sides before cancelling them. Don't just guess which ions are involved.
Using the wrong test reagent: Always add dilute nitric acid before testing for halides with silver nitrate, or dilute hydrochloric acid before testing for sulfates with barium chloride. This removes interfering ions.
Predicting reactions without checking the reactivity series: Always check the series before predicting whether a displacement reaction will occur. Don't assume based on what "seems right".
Exam technique for "C4: Predicting and Identifying Reactions and Products"
"Predict" questions require you to state what will happen AND provide reasoning. Simply writing an equation without explanation typically earns only half marks. Use phrases like "because X is more reactive than Y" or "because the product is insoluble".
Test identification questions follow a pattern: state the observation (colour change, precipitate), state what the observation indicates (which ion), then write the conclusion. Each step typically earns separate marks.
Symbol equation questions require balanced equations with correct formulas. Check your charges: sulfate is SO₄²⁻ not SO₄⁻. State symbols may be required for full marks—read the question carefully.
"Explain using ionic equations" means write the net ionic equation showing only participating ions, not the full equation with spectator ions. This is worth checking if the question specifically asks for it.
Quick revision summary
Chemical reactions follow predictable patterns based on reactivity series, compound types, and functional groups. Use the reactivity series to predict displacement, metal extraction, and reactions with acids or water. Acids produce specific salts depending on the acid type and base used. Precipitation reactions can be predicted using solubility rules. Identify ions using flame tests (cations) and chemical tests with specific reagents (both cations and anions). Write ionic equations by removing spectator ions from full ionic equations. Always justify predictions with scientific reasoning referencing reactivity or chemical properties.