WJEC GCSE Chemistry
WJEC GCSE Chemistry 1.6 Limestone Exam Questions
Free WJEC GCSE Chemistry revision for 1.6 Limestone: what you need to know, common mistakes and exam-style questions with answers. Covers thermal decomposition of carbonates, limestone, quicklime and slaked lime, the limestone cycle, uses of limestone and quarrying impacts.
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About WJEC GCSE Chemistry 1.6 Limestone
Limestone completes Unit 1 of WJEC GCSE Chemistry and is a good test of equation writing. You need to explain the trend in thermal stability of metal carbonates and write equations for their decomposition into oxides and carbon dioxide. You must know calcium carbonate, calcium oxide and calcium hydroxide as limestone, quicklime and slaked lime, and describe the cycle of reactions that links them, including limewater and carbon dioxide. Uses of limestone in iron and steel making, roads, neutralising soil and cement are covered, as are the impacts of quarrying. The specified practical compares the thermal stability of calcium, copper(II) and sodium carbonates.
What you need to know
- Thermal decomposition Metal carbonates break down on heating to a metal oxide and carbon dioxide, e.g. CaCO3 → CaO + CO2. Carbonates of more reactive metals are more stable.
- Specified practical Heat calcium carbonate, copper(II) carbonate and sodium carbonate and pass the gas through limewater. Copper(II) carbonate decomposes easily (green to black), calcium carbonate needs strong heating, and sodium carbonate does not decompose with a Bunsen burner.
- Names Limestone is calcium carbonate (CaCO3), quicklime is calcium oxide (CaO) and slaked lime is calcium hydroxide (Ca(OH)2). A solution of calcium hydroxide is limewater.
- Limestone cycle CaO + H2O → Ca(OH)2 is very exothermic. Ca(OH)2 + CO2 → CaCO3 + H2O turns limewater milky because insoluble calcium carbonate forms.
- Uses Limestone is used in the blast furnace to remove impurities in iron and steel making, for road building, in cement production, and (as limestone or lime) to neutralise acidic soil and lakes.
- Quarrying Benefits include jobs, income and useful building materials. Drawbacks include noise, dust, heavy traffic, loss of habitats and visual impact. Be ready to weigh these up in an evaluation.
Common mistakes in 1.6 Limestone
- Writing CaCO3 → Ca + CO3 for thermal decomposition. The products are calcium oxide and carbon dioxide: CaCO3 → CaO + CO2.
- Confusing quicklime and slaked lime. Quicklime is calcium oxide; slaked lime is calcium hydroxide, made by adding water to quicklime.
- Saying limewater turns white. It turns milky or cloudy because a white precipitate of calcium carbonate forms.
- Giving only drawbacks when asked to evaluate quarrying. Give benefits and drawbacks for local people, the economy and the environment, then a conclusion.
Try three 1.6 Limestone exam-style questions
1. Write a balanced equation for the reaction of calcium oxide with water. (1 mark)
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CaO + H2O → Ca(OH)2.
2. Calculate the mass of calcium oxide that can be made from 200 tonnes of calcium carbonate. (Mr CaCO3 = 100, Mr CaO = 56) (2 marks)
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100 tonnes gives 56 tonnes; 200 tonnes gives 112 tonnes.
3. Describe how a student could compare the thermal stability of three metal carbonates. (4 marks)
Show the mark scheme answer
Heat equal masses of each carbonate in a test tube with a Bunsen burner; pass gas through limewater; time how long until limewater turns milky; shortest time is least stable (keep heating strength the same).
1.6 Limestone FAQs
What does WJEC GCSE Chemistry 1.6 Limestone cover?
It covers the thermal decomposition and stability of metal carbonates, limestone, quicklime and slaked lime, the limestone cycle and limewater, uses of limestone and the impacts of quarrying.
Which unit is Limestone examined in?
It is examined in Unit 1: Chemical Substances, Reactions and Essential Resources, a 1 hour 45 minute written paper worth 45% of the GCSE, sat at foundation or higher tier. The Unit 3 practical assessment draws on the same practical skills.
Which equations should I learn for limestone?
Learn CaCO3 → CaO + CO2, CaO + H2O → Ca(OH)2 and Ca(OH)2 + CO2 → CaCO3 + H2O, plus the decomposition of copper(II) carbonate: CuCO3 → CuO + CO2.
