Exothermic and Endothermic ReactionsAQA GCSE Chemistry: Revision notes
Section 1
What are exothermic reactions?
Exothermic reactions are reactions that transfer energy to the surroundings, usually in the form of heat. This causes the temperature of the surroundings to rise.
In exothermic reactions, the enthalpy change (ΔH) is negative, indicating that energy is released overall.
Common examples of exothermic reactions include:
- Combustion – burning of fuels such as natural gas, petrol, or wood
- Neutralisation – reactions between acids and bases (e.g. HCl + NaOH)
- Oxidation reactions – such as rusting of iron or respiration
- Condensation of gases
- Freezing of liquids
- Many synthesis reactions
You can detect exothermic reactions practically because:
- A thermometer placed in the reaction mixture shows an increase in temperature
- The beaker or container becomes warm to touch
Examiners expect you to link the negative ΔH value directly to energy being released. Always state 'energy is transferred to the surroundings' rather than vague statements like 'energy is released'.
Think of exothermic reactions like spending money from your account – the system loses energy (negative change), and the surroundings gain it (the money goes elsewhere).
Section 2
What are endothermic reactions?
Endothermic reactions are reactions that take in energy from the surroundings. This causes the temperature of the surroundings to fall.
In endothermic reactions, the enthalpy change (ΔH) is positive, indicating that energy is absorbed overall.
Common examples of endothermic reactions include:
- Thermal decomposition – breaking down compounds using heat (e.g. calcium carbonate → calcium oxide + carbon dioxide)
- Reaction of citric acid with sodium hydrogen carbonate – produces a temperature drop
- Melting of solids
- Evaporation of liquids
- Photosynthesis
- Some dissolving processes (e.g. dissolving ammonium nitrate in water)
You can detect endothermic reactions practically because:
- A thermometer placed in the reaction mixture shows a decrease in temperature
- The beaker or container becomes cold to touch
- In extreme cases, water can freeze on the outside of a container
When describing endothermic reactions, examiners want to see 'energy is taken from the surroundings' and a clear connection to the positive ΔH value. Avoid saying 'the reaction needs heat' without explaining where that heat comes from.
Endothermic reactions are like withdrawing money from your account – the system gains energy (positive change), taking it from the surroundings (which now have less).
Section 3
How do energy level diagrams show reaction pathways?
Energy level diagrams (also called reaction profiles or energy profiles) show the energy changes that occur during a chemical reaction.
Key features of energy level diagrams:
- The vertical axis represents energy (usually in kJ/mol)
- The horizontal axis represents the progress of the reaction (reaction pathway)
- The diagram shows the energy of reactants at the start and the energy of products at the end
- The activation energy (Ea) is shown as the energy barrier that must be overcome
For exothermic reactions:
- Reactants start at a higher energy level than products
- The energy difference between reactants and products is shown as a negative value (ΔH)
- The overall energy 'goes down' on the diagram
- The activation energy is still required to start the reaction
For endothermic reactions:
- Reactants start at a lower energy level than products
- The energy difference between reactants and products is shown as a positive value (ΔH)
- The overall energy 'goes up' on the diagram
- The activation energy is still required to start the reaction
Both types of reactions require activation energy – this is the minimum energy needed for reactant particles to collide and react successfully. Without this energy barrier being exceeded, the reaction cannot occur, regardless of whether it is exothermic or endothermic overall.
When interpreting diagrams, examiners expect you to identify the activation energy as the difference between the reactants' energy level and the peak, and ΔH as the difference between the final products and initial reactants. Always label these correctly.
A common error is confusing activation energy with ΔH. Remember: activation energy is always positive and is needed to start any reaction; ΔH can be negative (exothermic) or positive (endothermic) and shows the overall energy change.
In an exothermic combustion reaction, if reactants have energy of 500 kJ/mol and products have energy of 300 kJ/mol, then ΔH = 300 − 500 = −200 kJ/mol. If the peak (transition state) is at 800 kJ/mol, then activation energy = 800 − 500 = 300 kJ/mol.
Section 4
What is activation energy and why does every reaction need it?
Activation energy (Ea) is the minimum energy required for a reaction to occur. It represents the energy needed for reactant particles to collide with enough force and correct orientation to form new bonds.
Why activation energy is essential:
- Reactant particles must collide with sufficient energy to break existing bonds
- Without reaching the activation energy threshold, collisions will simply bounce off without reaction occurring
- Even exothermic reactions (which release energy overall) still require activation energy to get started
- Only when particles have at least this minimum energy can the reaction proceed
Key points about activation energy:
- It is always a positive value, even for exothermic reactions
- It is the difference between the energy of the reactants and the peak (transition state) on an energy diagram
- Different reactions have different activation energies – some reactions are easier to start than others
- A reaction with high activation energy will be slow at room temperature
- A reaction with low activation energy will be fast at room temperature
- Catalysts work by lowering the activation energy, allowing reactions to proceed more quickly without being consumed themselves
Examiners often ask why a reaction is exothermic overall but still needs a spark or heat to start. The answer is always: 'activation energy must be provided first, then energy is released during the reaction'.
Combustion of natural gas is highly exothermic (ΔH very negative), but a match or spark is needed to start it. The spark provides the activation energy; once started, the exothermic reaction releases enough energy to continue.
Section 5
How do you compare exothermic and endothermic reactions?
| Feature | Exothermic | Endothermic |
|---|---|---|
| Energy change | Energy released to surroundings | Energy absorbed from surroundings |
| Temperature change | Temperature of surroundings increases | Temperature of surroundings decreases |
| Enthalpy change (ΔH) | Negative (ΔH < 0) | Positive (ΔH > 0) |
| Energy diagram | Products lower than reactants | Products higher than reactants |
| Activation energy | Always required (positive value) | Always required (positive value) |
| Examples | Combustion, neutralisation, rusting | Thermal decomposition, melting, photosynthesis |
| Practical observation | Beaker gets hot | Beaker gets cold |
| Everyday examples | Burning wood, neutralising acid | Ice melting (needs heat input), dissolving ammonium nitrate |
Important: Both exothermic and endothermic reactions require activation energy. The difference is only in whether energy is released overall (exothermic) or absorbed overall (endothermic).
In exam questions asking you to compare, use the table structure: state whether ΔH is positive or negative, describe the temperature change, give the direction on an energy diagram, and provide a relevant example. This systematic approach covers all marks.
Must Know
- Exothermic reactions release energy to the surroundings (ΔH negative), causing temperature to rise. Examples: combustion, neutralisation, oxidation reactions.
- Endothermic reactions absorb energy from the surroundings (ΔH positive), causing temperature to fall. Examples: thermal decomposition, reaction of citric acid with sodium hydrogen carbonate.
- Energy level diagrams show reactant and product energy levels. For exothermic reactions, products are lower in energy; for endothermic reactions, products are higher in energy.
- Activation energy (Ea) is the minimum energy required for any reaction to occur. It is always positive and is shown as the difference between reactants' energy and the peak on an energy diagram.
- Both exothermic and endothermic reactions require activation energy to start, even though exothermic reactions release energy overall.
- ΔH (enthalpy change) and activation energy are different: ΔH shows the overall energy change (negative or positive), while Ea is always positive and is needed to begin the reaction.
That's the notes covered.
Carry on to the next subtopic.