Calorimetry — AQA A-Level Chemistry
Test yourself on Calorimetry with AQA A-Level practice questions.
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Calorimetry explained
In a calorimetry experiment the reaction does not cool or warm itself: it warms or cools the solution around it, and q = mc∆T measures the heat gained or lost by that solution.
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The m is the mass in grams of the liquid whose temperature you measured, usually taken as the volume in cm³ because the solution is assumed to have the density of water, 1.00 g cm⁻³. The c is the specific heat capacity, the energy needed to raise one gram by one kelvin, 4.18 J g⁻¹ K⁻¹ for water. Because a kelvin and a degree Celsius are the same size, ∆T can be read straight off the thermometer. The answer comes out in joules. A temperature rise means the surroundings gained heat, so the reaction released it and ∆H is negative.
Students should be able to: use this equation to calculate the molar enthalpy change for a reaction use this equation in related calculations.
A molar enthalpy change needs two numbers: the heat, from q = mc∆T, and the amount in moles of the substance the question asks you to quote it for. Work out q in joules, divide by 1000, then divide by the moles of the limiting reactant, and attach the sign. So dissolving 4.00 g of ammonium nitrate, 0.0500 mol, in 50.0 cm³ of water and recording a fall of 6.2 K gives q = 50.0 × 4.18 × 6.2 = 1296 J, so ∆H = +1.296 ÷ 0.0500 = +25.9 kJ mol⁻¹, positive because the solution cooled. The same equation runs backwards: given ∆H and a temperature change you can find the mass or the concentration, and given the power of a heater, energy is power × time instead of mc∆T.
Required practical 2 Measurement of an enthalpy change.
Required practical 2 measures an enthalpy change directly. A known volume of solution goes into an insulated polystyrene cup with a lid. The starting temperature is recorded for a few minutes, the second reagent is added, and the temperature is followed until it falls steadily. Plotting temperature against time and extrapolating the cooling line back to the moment of mixing gives the temperature change with no heat loss, the ∆T for q = mc∆T. For example, adding zinc to copper sulfate solution. Design decisions are defensible: a lid and insulation cut heat exchange, stirring gives a uniform temperature, and an excess of one reagent guarantees the other reacts completely.
Your focus
- Identify which substance's mass and specific heat capacity belong in q = mc∆T for a dissolving, neutralisation or displacement experiment.
- Calculate q in joules from a recorded volume, temperature change and given specific heat capacity, then express it in kilojoules.
- Justify the sign given to ∆H by stating the direction of heat transfer between system and surroundings.
Show all 11 objectives
- Suggest a change of method that increases ∆T and explain why this lowers the percentage uncertainty in the temperature reading.
- Calculate a molar enthalpy change from raw calorimetry data by combining q = mc∆T with the amount in moles of the limiting reactant.
- Rearrange q = mc∆T to find an unknown mass, temperature change or specific heat capacity from a quoted enthalpy change.
- Calculate an enthalpy of vaporisation from the electrical energy supplied by a heater of stated power and the mass of liquid lost.
- Identify the limiting reactant in a calorimetry experiment and justify the choice before dividing.
- Describe a procedure that measures the enthalpy change of a reaction in an insulated cup.
- Determine a corrected temperature change by extrapolating the cooling line of a temperature against time graph back to the time of mixing.
- Calculate the percentage uncertainty in ∆T from the resolution of the thermometer.
Calorimetry exam tips
Marking Points
- one mark for using the mass of the solution, not the mass of the solid added, as m in q = mc∆T
- one mark for correct substitution into q = mc∆T with matching units, giving q in joules
- one mark for converting joules to kilojoules before dividing by the number of moles
- one mark for assigning the correct sign, negative ∆H for a temperature rise and positive ∆H for a temperature fall
- one mark for the amount in moles of the substance the enthalpy change is to be quoted per mole of
- one mark for q calculated correctly from q = mc∆T, or from power × time when a heater is used
- one mark for dividing the energy in kilojoules by the moles to reach kJ mol⁻¹
- one mark for the final value with the correct sign and units, allowing error carried forward from an earlier slip
- One mark for recording temperature before, during and after mixing and plotting it against time.
- One mark for extrapolating the cooling section back to the time of mixing to obtain a corrected ∆T.
- One mark for identifying a design feature that reduces heat loss, such as a lid, insulation or a polystyrene cup.
- One mark for using an excess of one reagent so the other is fully reacted, justified by which amount is used in the final division.
Examiner Tips
- 💡Write q = mc∆T, then label m, c and ∆T with the number and unit you are putting in, before you press a key.
- 💡Divide by 1000 once and say so in your working, so an examiner can follow the joule to kilojoule step.
- 💡If a question asks how to reduce the percentage uncertainty in ∆T, increase the temperature change, for example by using a larger mass of solid in the same volume of water.
- 💡Set the calculation out in three labelled lines, q, n, then ∆H = −q/n, so process marks survive an arithmetic slip.
- 💡Quote the answer to the same number of significant figures as the least precise data, usually three.
- 💡When a heater is specified in kW, energy in kJ is power in kW multiplied by time in seconds.
- 💡The absolute uncertainty for a temperature change (two readings) equals the thermometer's resolution (e.g., ±1 °C for 1 °C divisions), not double the resolution.
- 💡To improve the experiment, change one variable and say what it does: more solid in the same volume of water increases ∆T and lowers the percentage uncertainty.
Common Mistakes
- using the mass of the dissolved solid or of the metal rather than the mass of the water that changed temperature
- treating 25.0 cm³ of solution as 25.0 kg, or leaving the answer in joules and calling it kJ mol⁻¹
- converting ∆T from °C to K by adding 273, when only the difference is needed and the size of the unit is identical
- forgetting the minus sign when the temperature rose, so an exothermic reaction is reported as endothermic
- using the specific heat capacity of the metal or solid when the temperature measured was that of the water
- dividing by the moles of the excess reagent instead of the limiting one
- dividing joules by moles and then quoting the answer as kJ mol⁻¹ without the factor of 1000
- losing the sign at the last step, so an exothermic result is written as positive
- rounding the moles to one significant figure early, which moves the final answer outside the accepted range
- using the mass of water only when the question tells you the total mass of the mixture changed temperature
- Reading the highest temperature reached rather than the extrapolated value, which underestimates the enthalpy change.
- Starting the timer at the moment of mixing so there is no pre-mixing line to extrapolate from.
- Dividing by the moles of the reagent in excess rather than the limiting one.