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4.1.5Additional lessons

Calorimetry Experiments — IB Chemistry SL Revision Notes & Topical Questions

5 min read 6 parts Lesson 50 of 98

📚 IB Chemistry SL Hub / Topical Revision
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Part 1 of 6The Principle

4.1.5 Calorimetry Experiments

Calorimetry measures energy changes in the laboratory from temperature measurements.

Energy transfer
  • The heat released or absorbed by a reaction is passed to a known mass of water, whose temperature change is measured.
Equation
  • The heat transferred is q = mcΔT, where m is the mass of water, c its specific heat capacity and ΔT the temperature change.
Water values
  • The specific heat capacity of water is 4.18 J g⁻¹ K⁻¹, and 1.00 cm³ of dilute solution is assumed to have a mass of 1.00 g.
Sign of ΔH
  • Divide q by the moles reacting and change the sign if the temperature rose, because heat gained by water was lost by the reaction.

Part 2 of 6Typical Experiments

Typical Experiments
Solution reactions
  • Reactants are mixed in an insulated polystyrene cup and the maximum or minimum temperature is recorded as the reaction takes place.
Combustion
  • A spirit burner warms water in a metal can above it, and the burner is weighed before and after to find the mass of fuel burnt.
Measurements
  • Mass or volume of water, initial and final temperature, and the mass or moles of the limiting reactant are required.
Insulation
  • A lid and an insulating cup reduce the loss of heat to the surroundings, giving a more accurate temperature change.

Part 3 of 6Method for the Calculation

1
Find the mass of water or solution that changed temperature and calculate ΔT from the highest or lowest temperature reached.
2
Calculate q = mcΔT in joules and convert to kilojoules by dividing by 1000.
3
Work out the amount in moles of the limiting reactant using n = cV for solutions or n = m/M for solids.
4
Divide q by n to obtain the enthalpy change per mole and give ΔH a negative sign for a temperature rise or a positive sign for a fall.
5
State the answer to the same number of significant figures as the least precise measurement.

Part 4 of 6Errors and Improvements

Errors and Improvements
Heat loss
  • Heat escapes to the surroundings,
    • so the measured temperature rise is too small and the calculated ΔH is less exothermic than the true value.
Incomplete combustion
  • Soot formation in fuel experiments means less energy is released than for complete combustion,
    • so the result is too low in magnitude.
Assumptions
  • Treating the solution as water, ignoring the heat capacity of the cup, and assuming a density of 1.00 g cm⁻³ all introduce error.
Improvements
  • Use a lid, add insulation, stir continuously,
    • and plot a temperature–time graph extrapolated back to the time of mixing to correct for cooling.

Part 5 of 6Worked Example

Question

  • 50.0 cm³ of 1.00 mol dm⁻³ HCl is mixed with 50.0 cm³ of 1.00 mol dm⁻³ NaOH.
  • The temperature rises by 6.8 K.
  • Calculate ΔH for the neutralisation per mole of water formed.

  1. Mass of solution = 100.0 g, so q = 100.0 × 4.18 × 6.8 = 2842 J = 2.84 kJ.
  2. Amount of HCl = 1.00 × 0.0500 = 0.0500 mol, which equals the moles of water formed.
  3. Energy per mole = 2.84 ÷ 0.0500 = 56.8 kJ mol⁻¹.
  4. The temperature rose, so the reaction is exothermic and ΔH is negative.

Answer ΔH = −56.8 kJ mol⁻¹

Part 6 of 6Exam Checklist & Tips

4.1.5 – Key Points for the Exam

  • Calorimeter An insulated container in which the temperature change of a reaction is measured.
  • Specific heat capacity
    • The energy needed to raise the temperature of 1 g of a substance by 1 K;
      • for water it is 4.18 J g⁻¹ K⁻¹.
  • q = mcΔT The equation that gives the heat transferred from the mass, specific heat capacity and temperature change.
  • Limiting reactant The reactant that is completely used up, which determines the amount of heat released.

Common Mistakes

  • Using the mass of the reactants only, rather than the total mass of solution that is actually heated.
  • Leaving q in joules while ΔH is expected in kJ mol⁻¹.
  • Dividing by the moles of the excess reactant instead of the limiting reactant.
  • Forgetting the sign change, giving a positive ΔH for a reaction in which the temperature rises.

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