Calorimetry Calculations: q=mcΔT and Molar Enthalpy Change Revision Notes | A Level AQA Chemistry
Calorimetry uses q = mcΔT: heat change = mass × specific heat capacity × temperature change. Free AQA notes with 2 worked examples and 14 questions answered.
Updated 5 Sept 2026 · Revise Right Team
Calorimetry Calculations: q=mcΔT and Molar Enthalpy Change is part of Energetics in A Level AQA Chemistry.
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What this lesson covers
Calorimetry uses q = mcΔT: heat change = mass × specific heat capacity × temperature change. Free AQA notes with 2 worked examples and 14 questions answered.
Apply the equation q = mcΔT to calculate the heat change in a calorimetry experiment
Calculate the molar enthalpy change for a reaction using q = mcΔT and the number of moles
Explain the assumptions made in calorimetry calculations and evaluate their effect on accuracy
Analyse sources of error in calorimetry experiments and suggest improvements
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Use this lesson to stay inside Energetics while connecting the detail back to the wider Physical chemistry topic.
Heat change (q): The quantity of thermal energy transferred during a chemical reaction, measured in joules (J). Calculated using q = mcΔT and used to determine enthalpy changes.
Specific heat capacity (c): The energy required to raise the temperature of 1 gram of a substance by 1 °C, measured in J g⁻¹ °C⁻¹. For water, c = 4.18 J g⁻¹ °C⁻¹; this value is provided in the exam.
Molar enthalpy change (ΔH): The enthalpy change per mole of a specified reactant or product, measured in kJ mol⁻¹. Calculated by dividing q (in kJ) by the number of moles of the limiting reactant.
Temperature change (ΔT): The difference between the final and initial temperature of the solution or substance in a calorimetry experiment, measured in °C. ΔT = T_final − T_initial.
Revision notes
Read the core explanations from Calorimetry Calculations: q=mcΔT and Molar Enthalpy Change before testing yourself from memory.
1. Why Calorimetry Matters
Every time a fuel burns or a salt dissolves, energy moves — but how do chemists put a number on that? Calorimetry is the experimental technique that captures heat changes and converts them into the precise enthalpy values you see in data books.
At A Level, you need to do two things: use the equation q = mcΔT to find the heat change in joules, then scale that up to a molar enthalpy change in kJ mol⁻¹. These two steps are the backbone of every energetics calculation question on the AQA paper.
2. The Heat Equation: q = mcΔT
Picture a polystyrene cup containing 50 cm³ of water. A neutralisation reaction occurs inside it, and the temperature rises from 21.0 °C to 27.4 °C. The water has absorbed that energy — and q = mcΔT is the tool that tells you exactly how much.
The mass m refers to the substance that actually changes temperature — in most A Level calorimetry questions, this is the aqueous solution in the cup. Because dilute aqueous solutions are mostly water, the standard assumption is that the solution has the same specific heat capacity as water: c = 4.18 J g⁻¹ °C⁻¹. You will always be given this value; you are not expected to recall it.
One subtlety worth noting: ΔT is always T_final minus T_initial. For an exothermic reaction the temperature rises, so ΔT is positive and q comes out positive. The sign of ΔH is then assigned separately using the convention that exothermic reactions have a negative ΔH.
3. From Heat Change to Molar Enthalpy Change
Calculating q gives you the heat transferred for the specific quantities used in your experiment. But chemists need a value that is independent of scale — one that applies to exactly one mole of reactant. That is the molar enthalpy change, ΔH.
The route from q to ΔH has two steps: convert q from joules to kilojoules, then divide by the number of moles of the limiting reactant. The result carries units of kJ mol⁻¹ and a sign that reflects whether the reaction is exothermic (negative) or endothermic (positive).
The negative sign in ΔH = −q/n is a common source of confusion. Think of it this way: when the solution warms up, the surroundings have gained energy, meaning the reaction system has lost energy — so ΔH for the reaction is negative. The equation encodes that convention automatically.
4. Worked Example: Neutralisation Reaction
Theory becomes clear through a concrete calculation. The following example walks through a standard AQA-style neutralisation question, showing every step from raw data to a final molar enthalpy change.
Worked example: Calculating Molar Enthalpy Change for Neutralisation
25.0 cm³ of 1.00 mol dm⁻³ hydrochloric acid is added to 25.0 cm³ of 1.00 mol dm⁻³ sodium hydroxide solution in a polystyrene cup. The temperature rises from 19.5 °C to 25.9 °C. Calculate the molar enthalpy change of neutralisation. (c = 4.18 J g⁻¹ °C⁻¹; assume the density of the solution is 1.00 g cm⁻³.)
$m = (25.0 + 25.0) \text{ cm}^3 \times 1.00 \text{ g cm}^{-3} = 50.0 \text{ g}$ — Both solutions are mixed together, so the total volume is 50.0 cm³. Using density = 1.00 g cm⁻³, mass = volume numerically. This is the mass that undergoes the temperature change.
$\Delta T = 25.9 - 19.5 = 6.4 \text{ °C}$ — ΔT = T_final − T_initial. The temperature rose, so ΔT is positive, confirming an exothermic process.
$q = 50.0 \times 4.18 \times 6.4 = 1337.6 \text{ J}$ — Substituting m = 50.0 g, c = 4.18 J g⁻¹ °C⁻¹, ΔT = 6.4 °C. Arithmetic check: 50.0 × 4.18 = 209.0; then 209.0 × 6.4 = 1337.6 J.
$q = \frac{1337.6}{1000} = 1.3376 \text{ kJ}$ — Molar enthalpy changes are reported in kJ mol⁻¹, so convert by dividing by 1000.
$n(\text{HCl}) = 1.00 \times \frac{25.0}{1000} = 0.0250 \text{ mol}$ — n = concentration × volume (in dm³). Both HCl and NaOH react in a 1:1 ratio and are present in equal moles, so either can be used. Using HCl: 1.00 mol dm⁻³ × 0.0250 dm³ = 0.0250 mol.
$\Delta H = \frac{-1.3376}{0.0250} = -53.5 \text{ kJ mol}^{-1}$ — Dividing −q by n gives the molar enthalpy change. The negative sign confirms the reaction is exothermic. Arithmetic check: 1.3376 ÷ 0.0250 = 53.504, rounded to 3 significant figures = 53.5 kJ mol⁻¹.
Answer: ΔH = −53.5 kJ mol⁻¹ (exothermic)
5. Worked Example: Combustion in a Calorimeter
Combustion experiments use a slightly different set-up: a spirit burner heats a known mass of water directly, and the mass of fuel burned is found by weighing before and after. The same equation applies — but now m is the mass of water in the calorimeter, and the number of moles is calculated from the mass of fuel burned.
Worked example: Calculating Molar Enthalpy of Combustion
A student burns ethanol (M_r = 46.0 g mol⁻¹) under a copper calorimeter containing 200 g of water. The temperature of the water rises from 18.0 °C to 34.5 °C. The mass of the spirit burner decreases by 0.920 g. Calculate the molar enthalpy of combustion of ethanol. (c = 4.18 J g⁻¹ °C⁻¹)
$\Delta T = 34.5 - 18.0 = 16.5 \text{ °C}$ — T_final − T_initial. The water temperature rose, confirming heat was transferred from the combustion reaction to the water.
$q = 200 \times 4.18 \times 16.5 = 13794 \text{ J}$ — m = 200 g (mass of water), c = 4.18 J g⁻¹ °C⁻¹, ΔT = 16.5 °C. Arithmetic: 200 × 4.18 = 836.0; then 836.0 × 16.5 = 13794 J.
$q = \frac{13794}{1000} = 13.794 \text{ kJ}$ — Divide by 1000 to convert joules to kilojoules before calculating ΔH in kJ mol⁻¹.
$n = \frac{0.920}{46.0} = 0.0200 \text{ mol}$ — n = mass ÷ molar mass. The mass burned is the decrease in mass of the spirit burner: 0.920 g. Arithmetic: 0.920 ÷ 46.0 = 0.02000 mol.
$\Delta H = \frac{-13.794}{0.0200} = -689.7 \text{ kJ mol}^{-1}$ — ΔH = −q ÷ n. Arithmetic: 13.794 ÷ 0.0200 = 689.7 kJ mol⁻¹. The negative sign confirms combustion is exothermic. Rounded to 3 significant figures: −690 kJ mol⁻¹.
The data book value for the standard enthalpy of combustion of ethanol is around −1367 kJ mol⁻¹ — roughly twice the experimental value above. That large discrepancy is not a calculation error; it reflects the significant heat losses inherent in a simple spirit burner experiment. Evaluating these losses is a key A Level skill.
6. The Required Practical: Measuring Enthalpy Changes
AQA requires you to carry out calorimetry experiments and evaluate them critically. The practical below covers the standard solution-based approach — the same method underpins neutralisation, displacement, and dissolving experiments.
7. Assumptions, Errors, and Evaluation
Every calorimetry calculation rests on assumptions. Identifying these assumptions — and explaining how they introduce error — is exactly what AQA 'evaluate' and 'suggest improvements' questions are testing.
The specific heat capacity of the solution equals that of water (4.18 J g⁻¹ °C⁻¹). In reality, dissolved solutes alter c slightly, so this assumption underestimates or overestimates q.
The density of the solution is 1.00 g cm⁻³, so mass (g) equals volume (cm³) numerically. This is valid for dilute solutions but less accurate for concentrated ones.
All heat released by the reaction is absorbed by the solution — none is lost to the surroundings or absorbed by the calorimeter itself. In practice, heat loss is the largest source of error.
The reaction goes to completion, so all of the limiting reactant is consumed. Incomplete combustion in spirit burner experiments means less heat is released than the theoretical maximum.
No heat is lost through evaporation of solvent during the experiment.
For combustion experiments specifically, heat loss to the surroundings is severe because the flame is open to the air. The experimental value is almost always significantly less exothermic than the data book value. A bomb calorimeter, which encloses the reaction completely, minimises this loss and gives values much closer to the standard enthalpy of combustion.
8. Exam Technique and Common Mistakes
Calorimetry questions follow a predictable structure on AQA papers. Marks are allocated for the formula, correct substitution with units, and the final answer with sign and units. Losing the sign or the units at the end is one of the most common ways students drop marks they have otherwise earned.
Calorimetry at A Level is ultimately a two-step process: use q = mcΔT to find the heat change for the quantities used, then divide by moles to obtain the molar enthalpy change. Mastering the sign convention, unit conversion, and evaluation of assumptions covers every spec requirement for this topic and prepares you for the full range of AQA question styles.
Sources and content standards
Published by Revise Right. Use the official specification reference to check the requirements for your course. Content updated: 2026-09-05.
Calorimetry is the experimental technique that captures heat changes and converts them into enthalpy values. The temperature change of a known mass of water or solution is measured, q = mcΔT gives the heat change in joules, and this is scaled up to a molar enthalpy change in kJ mol⁻¹.
What does q = mcΔT mean?
q = mcΔT is the calorimetry equation for heat change: q is the heat transferred in joules, m is the mass of the substance that changes temperature (usually the solution), c is the specific heat capacity, 4.18 J g⁻¹ °C⁻¹ for water, and ΔT is the temperature change, final minus initial temperature.
What is a bomb calorimeter?
A bomb calorimeter is a calorimeter that encloses the combustion reaction completely. This minimises the heat loss to the surroundings that happens with an open spirit burner flame, so it gives values much closer to the standard enthalpy of combustion found in data books.
What assumptions are made in calorimetry calculations?
Calorimetry calculations assume the solution has the specific heat capacity of water, 4.18 J g⁻¹ °C⁻¹, that its density is 1.00 g cm⁻³ so mass equals volume, that all heat is absorbed by the solution with none lost, that the reaction goes to completion, and that no heat is lost through evaporation of solvent.
How do you calculate molar enthalpy change from calorimetry?
To calculate molar enthalpy change, find q using q = mcΔT, convert q to kilojoules by dividing by 1000, then divide by the moles of the limiting reactant: ΔH = −q/n. For 25.0 cm³ each of 1.00 mol dm⁻³ HCl and NaOH warming by 6.4 °C, ΔH = −53.5 kJ mol⁻¹.
Why is the experimental enthalpy of combustion lower than the data book value?
Experimental enthalpies of combustion are less exothermic than data book values mainly because of heat loss to the surroundings from the open flame, plus incomplete combustion in the spirit burner. A spirit burner experiment gave −690 kJ mol⁻¹ for ethanol, while the data book value is around −1367 kJ mol⁻¹.
Why is there a negative sign in ΔH = −q/n?
The negative sign in ΔH = −q/n is there because when the solution warms up, the surroundings have gained energy, meaning the reaction system has lost energy. So for an exothermic reaction, q is positive but ΔH for the reaction is negative, and the equation encodes that convention automatically.
How do you calculate the enthalpy of neutralisation from a temperature rise?
The enthalpy of neutralisation is found from q = mcΔT, then ΔH = −q/n. For 25.0 cm³ of 1.00 mol dm⁻³ HCl with 25.0 cm³ of 1.00 mol dm⁻³ NaOH rising by 6.4 °C, q = 50.0 × 4.18 × 6.4 = 1337.6 J, n = 0.0250 mol, so ΔH = −53.5 kJ mol⁻¹.
How do you calculate the enthalpy of combustion of ethanol from a calorimetry experiment?
The enthalpy of combustion of ethanol is found from the temperature rise of the water and the mass of ethanol burned. Heating 200 g of water by 16.5 °C gives q = 200 × 4.18 × 16.5 = 13794 J; 0.920 g of ethanol is 0.0200 mol, so ΔH = −13.794 ÷ 0.0200 = −690 kJ mol⁻¹.
How do you work out ΔT in calorimetry?
ΔT in calorimetry is always the final temperature minus the initial temperature. For an exothermic reaction the temperature rises, so ΔT is positive and q comes out positive; the sign of ΔH is then assigned separately, using the convention that exothermic reactions have a negative ΔH.
Why is the density of the solution assumed to be 1.00 g cm⁻³?
The density of the solution is assumed to be 1.00 g cm⁻³ so that the mass in grams equals the volume in cm³ numerically, for example 50.0 cm³ of solution has a mass of 50.0 g. This is valid for dilute solutions but less accurate for concentrated ones.
How does incomplete combustion affect a calorimetry result?
Incomplete combustion in a spirit burner experiment means less heat is released than the theoretical maximum, so the measured enthalpy of combustion is less exothermic than the true value. Calorimetry calculations assume the reaction goes to completion, so all of the limiting reactant is consumed.
How do you find the mass of fuel burned in a combustion experiment?
The mass of fuel burned in a combustion experiment is found by weighing the spirit burner before and after burning; the decrease in mass is the mass burned. For example, a decrease of 0.920 g of ethanol, with a molar mass of 46.0 g mol⁻¹, is 0.920 ÷ 46.0 = 0.0200 mol.
Why must q be converted to kJ before calculating ΔH?
The value of q must be converted to kJ, by dividing by 1000, because molar enthalpy changes are reported in kJ mol⁻¹. For example, 1337.6 J becomes 1.3376 kJ, which is then divided by the moles of the limiting reactant to give ΔH.