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1.4.2Physical Chemistry

Calorimetry — Oxford AQA International A Level Chemistry Revision Notes & Topical Questions

5 min read 6 parts Lesson 23 of 63

📚 Oxford AQA International A Level Chemistry Hub / Topical Revision
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Part 1 of 6The Heat Equation

1.4.2 Calorimetry
q = mcΔT
  • The heat energy transferred, q, equals the mass of water heated, times its specific heat capacity, times the temperature rise.
Mass
  • Use the mass of water (or solution) that is heated, in grams, with 1 cm³ of water taken as 1 g.
Specific heat capacity, c
  • Water has c = 4.18 J g⁻¹ K⁻¹, the energy needed to raise 1 g by 1 K.
Temperature change
  • ΔT is the final minus the initial temperature, which has the same size in °C and in K.

Part 2 of 6From q to ΔH

1
Calculate q in joules using q = mcΔT and convert to kJ by dividing by 1000.
2
Find the moles of the reactant or fuel that caused the temperature change, using n = m/M.
3
Divide q in kJ by the moles, to get the energy per mole.
4
Assign the sign: the temperature of the water rose so the reaction is exothermic and ΔH is negative.
5
State the answer in kJ mol⁻¹ with an appropriate number of significant figures.

Part 3 of 6Combustion Calorimetry

Combustion Calorimetry
Apparatus
  • A spirit burner containing the fuel heats a metal can of water, with the burner weighed before and after.
Measurements
  • Record the initial and highest water temperatures and the mass of fuel burned.
Precautions
  • Use a draught shield, stir the water, and place the flame close to the base of the can.
Error
  • Much heat is lost to the surroundings and combustion may be incomplete, so the measured value is usually less exothermic than the data book.

Part 4 of 6Assumptions

Heat capacity
  • The solution is assumed to have the same specific heat capacity and density as water.
No heat loss
  • All the heat is assumed to pass to the water and none to the apparatus or surroundings.
Complete reaction
  • The fuel is assumed to burn completely, forming carbon dioxide and water only.
Effect
  • These assumptions mean the measured enthalpy change is usually smaller in magnitude than the true one.

Part 5 of 6Worked Example

Question

  • 0.230 g of ethanol (Mr = 46.0) is burned and heats 150 g of water from 20.0 °C to 28.0 °C.
  • Calculate the enthalpy change of combustion of ethanol from these results.

  1. q = mcΔT = 150 × 4.18 × 8.0 = 5016 J = 5.016 kJ.
    What the symbols mean: E, Q or q = energy transferred (J) · m = mass (kg; g in chemistry calorimetry) · c = specific heat capacity (J kg−1 °C−1; water 4.18 J g−1 °C−1) · ΔT or Δθ = temperature change (°C or K)
  2. n(ethanol) = 0.230 / 46.0 = 0.00500 mol.
  3. Energy per mole = 5.016 / 0.00500 = 1003 kJ mol⁻¹.
  4. The reaction is exothermic, so ΔH is negative.

Answer ΔcH ≈ −1.00 × 10³ kJ mol⁻¹ (less exothermic than the data book value because of heat loss).

Part 6 of 6Exam Checklist & Tips

1.4.2 – Key Points for the Exam

  • Specific heat capacity The energy required to raise the temperature of 1 g of a substance by 1 K.
  • Calorimetry The measurement of heat energy changes using the temperature change of a substance.
  • Spirit burner A small container of liquid fuel with a wick, used in combustion experiments.
  • Heat loss Energy transferred to the surroundings instead of the water.

Common Mistakes

  • Using the mass of the fuel rather than the mass of the water in q = mcΔT.
  • Forgetting to convert joules to kilojoules before giving ΔH in kJ mol⁻¹.
  • Giving a positive value for a combustion enthalpy that is exothermic.
  • Using the wrong number of moles because Mr of the fuel was miscalculated.

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