OCR GCSE Chemistry A: Energetics (J248)

Sub-topic C3.2 Energetics sits in Topic C3 Chemical reactions of the OCR GCSE (9–1) Chemistry A (Gateway Science) J248 specification (Version 4.0, August 2026), and this guide teaches all of it: learning outcomes C3.2a to C3.2d, the mathematical learning outcomes CM3.2i and CM3.2ii, and the practical suggestion linked to PAG C8. Outcomes C3.2a, C3.2b and C3.2c are assessed at both tiers. Outcome C3.2d, calculating energy changes from bond energies, is Higher Tier only. In the specification’s assessment overview, Topic C3 content belongs to Paper 1 for Foundation Tier candidates and Paper 3 for Higher Tier candidates.

Once you have worked through the guide, condense it with the Energetics revision notes and test yourself with the Energetics practice questions. The full course is on the OCR GCSE Chemistry hub, and you can tick off every statement on the printable J248 checklist. A short free 10-minute diagnostic will show which parts of chemistry need the most work.

What this sub-topic covers

Reference	Skill	Tier
C3.2a	Tell exothermic from endothermic reactions by the temperature change of the surroundings	Both tiers
C3.2b	Draw and label reaction profiles for both types, showing activation energy, energy change, reactants and products	Both tiers
C3.2c	Explain activation energy as the energy needed for a reaction to occur	Both tiers
C3.2d	Calculate the energy change of a reaction from bond breaking and bond making energies	Higher Tier only
CM3.2i	Interpret charts and graphs of reaction profiles (maths skill M4a)	Both tiers
CM3.2ii	Carry out the arithmetic for energy changes (maths skill M1a)	Both tiers
PAG C8	Measure the temperature change in reactions	Both tiers

Energy is transferred, not used up

Before the definitions, fix one idea. In a reaction, energy is transferred between the chemicals and their surroundings. It is not created, lost or used up. By “surroundings” chemists mean the solvent, the container, the thermometer and the air: anything other than the substances actually reacting.

• In an exothermic reaction, energy is transferred from the reacting chemicals to the surroundings. The surroundings warm up, so the temperature reading rises.
• In an endothermic reaction, energy is transferred from the surroundings to the reacting chemicals. The surroundings cool down, so the temperature reading falls.

The thermometer measures the surroundings, so a fall means energy has gone into the chemicals.

Exothermic examples: combustion, neutralisation, and many metal–acid reactions. Endothermic examples: thermal decomposition of a metal carbonate, and hydrogencarbonates reacting with citric or ethanoic acid.

Worked example: classifying from temperature data (C3.2a)

A student mixes reagents in a polystyrene cup and records the temperature before mixing and the highest or lowest temperature reached.

Reaction	Start (°C)	Final (°C)
A: magnesium ribbon + dilute sulfuric acid	18.5	31.0
B: potassium hydrogencarbonate + ethanoic acid	20.0	14.2
C: potassium hydroxide + nitric acid	19.0	25.5

Step 1: temperature change = final − start.

• A: 31.0 − 18.5 = +12.5 °C
• B: 14.2 − 20.0 = −5.8 °C
• C: 25.5 − 19.0 = +6.5 °C

Step 2: interpret the sign. A rise means the surroundings gained energy, so A and C are exothermic. A fall means the surroundings lost energy to the chemicals, so B is endothermic.

Measuring temperature change in practice (PAG C8)

The specification’s practical suggestion for C3.2a is measuring the temperature change in reactions, which falls within PAG C8. A reliable method looks like this:

1. Measure a fixed volume of the first solution with a measuring cylinder and pour it into a polystyrene cup standing in a beaker (the beaker stops it tipping).
2. Record the starting temperature with a thermometer or temperature probe.
3. Add the second reagent (a measured volume of solution, or a weighed mass of solid), put a lid on and stir.
4. Record the highest temperature (or lowest, for an endothermic reaction).
5. Repeat the whole experiment and work out a mean temperature change.

What the cup and cover are for. Expanded polystyrene traps air, so heat passes through it slowly, and covering the cup stops warm or cool air escaping from the surface. Without them, energy leaks to or from the room, and the measured change is smaller than it should be every time. That is a systematic error: it pushes all results the same way. Misreading the thermometer scale, or catching the peak a few seconds late, gives random errors that scatter results above and below the true value.

Handling repeat data. Suppose four repeats give temperature rises of 6.4, 6.9, 6.8 and 8.7 °C. The 8.7 °C result is far from the others, so treat it as anomalous and leave it out. Mean = (6.4 + 6.9 + 6.8) ÷ 3 = 20.1 ÷ 3 = 6.7 °C. The range of the three is 6.9 − 6.4 = 0.5 °C, so one simple estimate of the uncertainty is half the range, ±0.25 °C. Results that sit close together are precise and show good repeatability. If another group using the same method gets similar values, the results are reproducible.

Reaction profiles (C3.2b, CM3.2i)

A reaction profile is a graph of energy, plotted up the page, against how far the reaction has gone, plotted across. It shows four things you must label:

• reactants: a horizontal line at the start
• products: a horizontal line at the end
• activation energy: an arrow from the reactants line up to the top of the curve (the peak)
• energy change: an arrow from the reactants line to the products line

Here is a sketch of an exothermic profile:

energy
  |     ^      .---.
  |     |     /     \
  |  Ea |    /       \
  |     |   /         \
  | ----+--'           \
  | reactants |         \
  |           | energy   \
  |           v change    '-------
  |                       products
  +----------------------------------
             progress of reaction

On an exothermic profile the products line sits below the reactants line, so the energy change arrow points down.

For an endothermic reaction, draw the products line higher than the reactants. The energy change arrow points up, from reactants to products. The peak is still above both lines, and the activation energy arrow still starts at the reactants.

Worked example: reading a profile

A profile shows the reactants at 250 kJ/mol, the top of the curve at 330 kJ/mol and the products at 60 kJ/mol.

• Activation energy = peak − reactants = 330 − 250 = 80 kJ/mol
• Energy change = products − reactants = 60 − 250 = −190 kJ/mol
• Products sit below reactants, which makes this reaction exothermic. The negative sign says the same thing: the chemicals end up with less energy, and 190 kJ/mol has gone to the surroundings.

Measure both arrows from the reactants line. A common slip is measuring activation energy from the products line or from the axis.

Activation energy (C3.2c)

Activation energy is the minimum energy needed for a reaction to occur. Particles must collide with at least this much energy for bonds in the reactants to start breaking.

This explains why many exothermic reactions need a “start”. A candle, a gas cooker or a firework needs a flame or spark first. That initial energy supplies the activation energy. After that, the energy released keeps it going. So a burning candle is exothermic, even though you had to light it. The specification lists this exact confusion as a common misconception: needing heat to start a reaction does not make it endothermic. Whether a reaction is exothermic or endothermic depends only on the overall energy change, from reactants to products.

On a profile, a large activation energy shows as a tall peak above the reactants.

Bond energies (C3.2d, Higher Tier only)

The bond model explains where the energy goes:

• To break any bond, energy must be supplied, so bond breaking is an endothermic process.
• When any bond forms, energy is given out, so bond making is an exothermic process.

A frequent mistake is to say energy is released when bonds break. It is the other way round. Energy is released only when new bonds form.

Each bond has a bond energy: the energy needed to break one mole of that bond, in kJ/mol. The same amount is released when one mole of that bond forms. Every calculation here supplies the bond energies to use.

Energy change = total energy to break bonds in the reactants − total energy released making bonds in the products

• A negative answer means bond making gave out more than bond breaking took in: exothermic.
• A positive answer means bond breaking took in more than bond making gave out: endothermic.

Method in four steps

1. Draw (or picture) the displayed formula of every molecule, so you can see every bond.
2. List and add up the bonds broken, multiplying by the number in the balanced equation.
3. List and add up the bonds made in the same way.
4. Subtract: broken − made. Give the sign and the unit, kJ/mol.

Worked example 1: ethene and bromine

CH₂=CH₂ + Br–Br → CH₂Br–CH₂Br

Use these values (kJ/mol): C=C 612, Br–Br 193, C–C 348, C–Br 276. The four C–H bonds are unchanged, so you can leave them out of both totals.

• Bonds broken: C=C + Br–Br = 612 + 193 = 805 kJ
• Bonds made: C–C + 2 × C–Br = 348 + (2 × 276) = 348 + 552 = 900 kJ
• Energy change = 805 − 900 = −95 kJ/mol, exothermic

Worked example 2: nitrogen and oxygen

N≡N + O=O → 2 N=O

Use these values (kJ/mol): N≡N 945, O=O 498, N=O 630.

• Bonds broken: 945 + 498 = 1443 kJ
• Bonds made: 2 × 630 = 1260 kJ
• Energy change = 1443 − 1260 = +183 kJ/mol, endothermic

Breaking the very strong triple bond in nitrogen takes in more energy than forming the two new bonds gives out. That strong bond also means a high activation energy, which is one reason the nitrogen and oxygen in air do not react at room temperature. They combine only at very high temperatures, such as inside a car engine or in a lightning strike.

Worked example 3: burning propane

C₃H₈ + 5O₂ → 3CO₂ + 4H₂O

Use these values (kJ/mol): O–H 464, C=O 805, C–C 348, O=O 498, C–H 413.

Propane has 2 C–C bonds and 8 C–H bonds. Each CO₂ has 2 C=O bonds and each H₂O has 2 O–H bonds.

Bonds broken	Energy (kJ)	Bonds made	Energy (kJ)
2 × C–C	696	6 × C=O	4830
8 × C–H	3304	8 × O–H	3712
5 × O=O	2490
Total	6490	Total	8542

Energy change = 6490 − 8542 = −2052 kJ/mol, exothermic. The equation is per mole of propane, so this is the energy change when one mole of propane burns.

The value from bond energies is an estimate. Bond energies are averages taken over many different molecules, so a calculated answer is usually a little different from a measured one.

Maths skills (CM3.2i and CM3.2ii)

Reading a profile (M4a) means taking levels off the energy axis and subtracting from the reactants level. Energy change arithmetic (M1a) means multiplying each bond energy by its number of bonds before adding, then keeping the sign.

Common errors

• Saying a temperature fall means the reaction “lost heat”. In an endothermic reaction the chemicals gain energy; the surroundings lose it.
• Drawing the activation energy arrow from the axis or from the products.
• Drawing an endothermic profile with the products below the reactants.
• Forgetting the multiplier: 5O₂ means five O=O bonds, not one.
• Missing a bond: counting one C=O for CO₂ instead of two.
• Writing “products − reactants” for bond energies. For bond energies it is broken − made.
• Dropping the sign on the final answer.

Next steps

Use the Energetics revision notes for a summary and self-test, then the Energetics practice questions. Relative formula mass, needed when energy changes are compared per gram, is in the purity and separating mixtures guide. Activation energy returns in the controlling reactions guide, which covers collision theory and catalysts.

Official syllabus

OCR GCSE (9–1) Chemistry A (Gateway Science) J248 specification, Version 4.0 (August 2026), for first assessment in 2018, published by OCR. Sub-topic C3.2 Energetics, learning outcomes C3.2a–C3.2d, mathematical learning outcomes CM3.2i–CM3.2ii and practical suggestions (PAG C8).