Calcium Carbonate

Calcium carbonate (CaCO3) is one of the most frequently examined compounds in 0620, running through acid reactions, thermal decomposition, the limewater test, rate-of-reaction experiments and industrial chemistry.

Key facts

Property	Value
Formula	CaCO3
Mr	40 + 12 + 3(16) = 100
Bonding	Ionic — Ca2+ and CO3^2-, giant ionic lattice
Key property	White solid, insoluble in water, decomposes on strong heating
Natural forms	Limestone, marble, chalk (same formula, different rock)
Main uses	Making cement and lime; neutralising acidic soil; blast furnace

Where calcium carbonate appears in 0620

• Rate of reaction: marble chips + acid is the standard experiment (varying surface area, concentration, temperature).
• Preparation of salts: a carbonate reacting with acid.
• Extraction of metals: limestone in the blast furnace.
• Air quality and climate change: acid rain attacking limestone buildings and statues.

Natural forms

The same compound occurs as three rocks with different physical properties set by how they formed:

• Limestone: sedimentary rock, quarried for building and manufacturing.
• Marble: metamorphic rock, formed from limestone under heat and pressure; harder, and takes a polish.
• Chalk: soft, porous sedimentary rock.

Key reactions

Thermal decomposition

CaCO3(s) -> CaO(s) + CO2(g)

An endothermic thermal decomposition needing strong heat. The product CaO (quicklime) is a white solid.

With acids

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

Observations: effervescence of CO2, the solid dissolves, and the gas turns limewater milky.

The limewater test (the reverse reaction)

Ca(OH)2(aq) + CO2(g) -> CaCO3(s) + H2O(l)

The white precipitate makes limewater cloudy — the standard test for carbon dioxide.

The limestone cycle

1. CaCO3(s) -> CaO(s) + CO2(g) (heat strongly)
2. CaO(s) + H2O(l) -> Ca(OH)2(s) (add water: vigorous, exothermic — slaking)
3. Ca(OH)2(aq) + CO2(g) -> CaCO3(s) + H2O(l) (bubble CO2 through limewater)

The cycle returns to the starting compound. See calcium oxide and calcium hydroxide.

Uses

Use	Explanation
Cement and concrete	Heated with clay to make cement; mixed with sand and gravel for concrete
Blast furnace	Decomposes to CaO, which removes acidic SiO2 as molten slag
Neutralising acidic soil	Powdered limestone raises soil pH without over-liming
Glass making	Heated with sand and soda

Marble chips react with dilute hydrochloric acid. The experiment is repeated using the same mass of marble ground into a fine powder. Explain, using collision theory, why the powder reacts faster. (3 marks)

Mark scheme

• The powder has a much larger (total) surface area than the chips [1]
• so more of the marble is exposed and there are more frequent collisions between acid particles and the surface [1]
• the frequency of successful collisions per second increases, so the rate of reaction increases [1]

Examiner note: connect surface area to how often particles collide, not to collision energy. Breaking a solid up does not change the activation energy or the temperature — only the collision frequency.

Common exam mistakes

• Saying thermal decomposition is exothermic. It is endothermic; heat must be supplied continuously to keep it going.
• Confusing the lime compounds: quicklime is calcium oxide (CaO), slaked lime is calcium hydroxide (Ca(OH)2), and limewater is dilute aqueous calcium hydroxide.
• Getting the acid ratio wrong. Calcium is 2+, so CaCO3 + 2HCl -> CaCl2 + H2O + CO2, not CaCl.
• Thinking powder has a smaller surface area. For the same mass, powder or small chips have a larger total surface area, so they react faster.