Calorimetry is the experimental technique of measuring temperature changes to calculate the energy transferred in a chemical reaction. By measuring how much a known mass of water heats up or cools down, and using the equation q = mcΔT, chemists can determine whether a reaction is exothermic or endothermic and calculate the enthalpy change per mole.
What is calorimetry and when is it used?
A calorimeter is any apparatus designed to measure heat transfer during a chemical or physical change. At GCSE, calorimetry is most commonly applied to:
- Combustion reactions (burning fuels or food) — the fuel heats water in a copper calorimeter or beaker.
- Neutralisation reactions — the temperature rise when an acid and alkali are mixed.
- Dissolving (enthalpy of solution) — the temperature change when a solid dissolves in water.
- Displacement reactions — temperature change when a more reactive metal displaces a less reactive one from a solution.
In every case the principle is the same: the energy released (or absorbed) by the reaction is transferred to (or from) a known mass of water, and the temperature change tells you how much energy moved.
What is the key equation for calorimetry?
The heat transferred to or from the water is calculated using:
q = m × c × ΔT
Where:
- q = heat energy transferred (joules, J)
- m = mass of water (grams, g)
- c = specific heat capacity of water = 4.18 J g⁻¹ °C⁻¹
- ΔT = temperature change (°C) = T_final − T_initial
A positive ΔT means the water warmed up → the reaction is exothermic. A negative ΔT means the water cooled down → the reaction is endothermic.
Worked example: neutralisation calorimetry
Question: 25.0 cm³ of 1.0 mol/dm³ hydrochloric acid is mixed with 25.0 cm³ of 1.0 mol/dm³ sodium hydroxide. The temperature rises from 21.0 °C to 27.6 °C. Calculate the enthalpy change of neutralisation in kJ/mol.
Step 1 — Mass of solution Assume the density of the dilute solution is 1.0 g/cm³. Total volume = 25.0 + 25.0 = 50.0 cm³ → mass = 50.0 g
Step 2 — Temperature change ΔT = 27.6 − 21.0 = 6.6 °C
Step 3 — Energy transferred q = m × c × ΔT = 50.0 × 4.18 × 6.6 = 1379 J = 1.38 kJ
Step 4 — Moles of water formed Moles of HCl = (25.0/1000) × 1.0 = 0.025 mol (Moles of NaOH = 0.025 mol — equal, so neutralisation is complete) Moles of water formed = 0.025 mol
Step 5 — Enthalpy change per mole ΔH = −q / moles = −1.38 / 0.025 = −55.2 kJ/mol
The negative sign shows the reaction is exothermic (energy released to surroundings). The literature value is −57.1 kJ/mol — the experimental value is lower because of heat losses.
What causes errors in simple calorimetry experiments?
Simple GCSE calorimetry setups have significant sources of error:
| Source of error | Effect | How to reduce it |
|---|---|---|
| Heat loss to the surroundings | Temperature rise is less than it should be → ΔH appears smaller | Use an insulated cup (polystyrene); add a lid; plot a temperature–time graph and extrapolate |
| Incomplete combustion | Less heat released than expected | Ensure adequate oxygen supply; clean wick |
| Specific heat capacity assumption | Using c = 4.18 for solution (which may not be exactly 4.18) | Use calibrated calorimeter; accept as a known approximation |
| Heat absorbed by apparatus | Thermometer and cup absorb heat, not counted in calculation | Use calorimeter calibration factor |
| Temperature not uniform | Thermometer reads one point in solution | Stir thoroughly before reading |
For combustion experiments, heat losses are particularly severe — burning a crisp or candle under a beaker is extremely inefficient, and measured values of ΔH are typically 30–60 % lower than literature values.
How do you convert a measured q to an enthalpy change?
Enthalpy change (ΔH) is defined per mole of a named substance. To convert:
- Calculate q using q = mcΔT.
- Calculate the moles of the substance of interest that reacted.
- ΔH (J/mol) = q ÷ moles
- Convert to kJ/mol by dividing by 1000.
- Assign sign: exothermic reactions have a negative ΔH (energy released); endothermic reactions have a positive ΔH (energy absorbed).
This sign convention is important: a temperature rise in the surroundings (water heats up) means the system (reaction) lost energy, so ΔH is negative.
How does a temperature–time graph improve accuracy?
In practice, a student records the temperature every 30 seconds before and after mixing reagents. This allows them to:
- Extrapolate the cooling curve back to the moment of mixing, finding the maximum temperature that would have been reached if there were no heat losses.
- Subtract the initial temperature from this extrapolated maximum to find a corrected ΔT.
Using the extrapolated ΔT instead of the observed maximum gives a more accurate enthalpy calculation and is the expected technique in higher-mark GCSE and A-level practical work.
Frequently asked questions
What equation do you use in calorimetry GCSE chemistry?
The key equation is q = mcΔT, where q is the heat energy in joules, m is the mass of water (or solution) in grams, c is the specific heat capacity of water (4.18 J g⁻¹ °C⁻¹), and ΔT is the temperature change in degrees Celsius. The sign of ΔT tells you whether the reaction is exothermic (ΔT positive, water warms up) or endothermic (ΔT negative, water cools down). You then divide q by the number of moles that reacted to get the enthalpy change per mole.
Why is the specific heat capacity of water used even when the solution is not pure water?
At GCSE, dilute aqueous solutions are assumed to have the same specific heat capacity and density as pure water (4.18 J g⁻¹ °C⁻¹ and 1.0 g/cm³). This is an approximation — the actual values vary with concentration — but it is acceptable for GCSE calculations and is a standard assumption stated in mark schemes. You should make this assumption explicit in your answer: "assuming c = 4.18 J g⁻¹ °C⁻¹ and density = 1.0 g/cm³."
Why do calorimetry experiments give values lower than the literature enthalpy change?
The main reason is heat loss to the surroundings. The water and solution transfer heat to the apparatus, the air, and any uninsulated surfaces, so the measured temperature rise is smaller than it would be in a perfectly insulated system. Other factors include incomplete reaction, heat absorbed by the calorimeter itself, and the approximation of the specific heat capacity. These systematic errors mean experimental ΔH values are almost always less negative (for exothermic reactions) than the true value.
What is the difference between q and ΔH in chemistry?
q is the heat energy measured in a specific experiment — it depends on the quantities of reactants used and the apparatus. ΔH is the standard enthalpy change, defined per mole of a specified substance under given conditions. To convert q to ΔH, you divide by the number of moles that reacted and apply the correct sign: ΔH is negative for exothermic reactions (heat is released to the surroundings) and positive for endothermic reactions (heat is absorbed from the surroundings).
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