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Calorimetry and enthalpy changes
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Measuring a temperature change lets you calculate the energy released or taken in by a reaction.
Exothermic and endothermic
Exothermic: heat energy is given out; the temperature of the surroundings rises. Examples: combustion, neutralisation, many displacement reactions.
Endothermic: heat energy is taken in; the temperature falls. Example: dissolving ammonium nitrate in water.
Endothermic: heat energy is taken in; the temperature falls. Example: dissolving ammonium nitrate in water.
Simple calorimetry
Solutions (neutralisation, displacement, dissolving): mix in an insulated polystyrene cup with a lid. Record the starting and highest (or lowest) temperature.
Combustion: burn a fuel in a spirit burner under a copper can of water. Weigh the burner before and after.
Q = m × c × ΔT
Q is heat energy (J), m is the mass of water or solution (g), c = 4.2 J/g °C, ΔT is the temperature change (°C). For solutions, take 1 cm3 as 1 g.
Combustion: burn a fuel in a spirit burner under a copper can of water. Weigh the burner before and after.
Q = m × c × ΔT
Q is heat energy (J), m is the mass of water or solution (g), c = 4.2 J/g °C, ΔT is the temperature change (°C). For solutions, take 1 cm3 as 1 g.
Molar enthalpy change
ΔH = −Q ÷ moles (convert Q to kJ first, giving kJ/mol).
ΔH is negative for exothermic reactions and positive for endothermic ones.
Example: 50 cm3 of acid + 50 cm3 of alkali (0.050 mol each) rise by 6.8 °C.
Q = 100 × 4.2 × 6.8 = 2856 J = 2.856 kJ. ΔH = −2.856 ÷ 0.050 = −57.1 kJ/mol
ΔH is negative for exothermic reactions and positive for endothermic ones.
Example: 50 cm3 of acid + 50 cm3 of alkali (0.050 mol each) rise by 6.8 °C.
Q = 100 × 4.2 × 6.8 = 2856 J = 2.856 kJ. ΔH = −2.856 ÷ 0.050 = −57.1 kJ/mol
Dissolving 0.10 mol of a salt in 50 g of water lowers the temperature by 5.0 °C. Find Q and ΔH.
- Q = 50 × 4.2 × 5.0 = 1050 J = 1.05 kJ
- The temperature fell, so it is endothermic and ΔH is positive.
- ΔH = +1.05 ÷ 0.10
Answer: Q = 1050 J; ΔH = +10.5 kJ/mol
Exothermic reactions raise the temperature; endothermic ones lower it. Q = m × c × ΔT with c = 4.2 J/g °C. ΔH = −Q ÷ moles, in kJ/mol: negative for exothermic, positive for endothermic.
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