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Learn5070Paper 1 · Multiple ChoiceChemical energeticsExothermic and endothermic reactions
Chemical energetics · Subtopic 1 of 1

Exothermic and Endothermic Reactions

Exothermic reactions release heat to the surroundings (temperature rises; ΔH\Delta H negative), e.g. combustion, neutralisation. Endothermic reactions take in heat (temperature falls; ΔH\Delta H positive), e.g. thermal decomposition, photosynthesis. Breaking bonds takes in energy; making bonds releases it. ΔH=∑(bonds broken)−∑(bonds made)\Delta H=\sum(\text{bonds broken})-\sum(\text{bonds made}). Energy released into water: Q=mcΔTQ=mc\Delta T.

Cambridge 5070 Paper 1Cambridge 5070 Paper 2Cambridge 0620 Papers 2, 4Cambridge 0620 Papers 1, 3Edexcel 4CH1 Paper 1CEdexcel 4CH1 Paper 2C≈ 45 minRevision 1 · updated 2026-10-30
Section 1 of 14

Learning objectives

  • Classify reactions as exothermic or endothermic.
  • Draw and interpret reaction pathway diagrams with activation energy.
  • Calculate ΔH\Delta H from bond energies.
  • Calculate energy changes with Q=mcΔTQ=mc\Delta T.
Pause and think

An instant cold pack gets icy when you snap it, with no freezer involved. Where does the "cold" come from?

Section 2 of 14

Core idea

Start hereHot and cold packs

Snap one pack and it heats up; snap another and it freezes. One reaction releases energy to the surroundings, the other absorbs it. Bonds explain both.

The big idea

Breaking bonds takes in energy; making bonds releases it. The overall energy change ΔH\Delta H is the difference: negative (exothermic) when more energy is released than absorbed.

1. Energy profiles

Reading a pathway diagram

The hump is the activation energy; the step is ΔH.

2. Bond energies

Where ΔH\Delta H comes from
  1. Ein=∑EbrokenE_{\text{in}}=\textstyle\sum E_{\text{broken}}
    Energy needed to break all reactant bonds.
  2. Eout=∑EmadeE_{\text{out}}=\textstyle\sum E_{\text{made}}
    Energy released forming all product bonds.
  3. ΔH=Ein−Eout\Delta H=E_{\text{in}}-E_{\text{out}}
    Net energy taken in by the system.
Worked exampleCombustion of methane

CH4+2O2→CO2+2H2O\mathrm{CH_4+2O_2\rightarrow CO_2+2H_2O}. C–H 413, O=O 498, C=O 805, O–H 464 kJ/mol.

Goal
ΔH\Delta H.
Strategy
Count bonds from displayed formulae.
  1. 4(413)+2(498)=26484(413)+2(498)=2648
    Broken: 4 C–H, 2 O=O.
  2. 2(805)+4(464)=34662(805)+4(464)=3466
    Made: 2 C=O, 4 O–H.
  3. ΔH=2648−3466=−818 kJ/mol\Delta H=2648-3466=-818\ \mathrm{kJ/mol}
    Broken − made.
Answer−818-818 kJ/mol (exothermic).
✓ CheckFuels are exothermic ✓.

3. Calorimetry

Q=mcΔTQ=mc\Delta T
where
Q
heat energy (J)
m
mass of water or solution (g)
c
4.2 J/g/°C
\Delta T
temperature change (°C)
Common misconceptionSign of ΔH If the temperature of the water rises, the reaction is exothermic, so ΔH=−Q/n\Delta H=-Q/n is negative.
Pause and think

Why do measured enthalpies of combustion usually come out smaller than data-book values?

Summary

  • Exo: ΔH<0\Delta H<0. Endo: ΔH>0\Delta H>0.
  • ΔH\Delta H = broken − made.
  • Q=mcΔTQ=mc\Delta T.
Key words
Exothermic
Transfers heat energy to the surroundings.
Endothermic
Takes in heat energy from the surroundings.
Activation energy
Minimum energy needed for a reaction to occur.
HistoryHess. In 1840 Germain Hess showed that the total energy change of a reaction is the same whatever route it takes, a form of energy conservation decades before it was generally accepted.
Fun factSelf-heating cans. Some coffee cans heat themselves: pressing a button mixes calcium oxide with water in a hidden compartment, and the exothermic reaction warms the drink.
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