Energetics of Reactions

Summary: Exothermic reactions release energy to surroundings (temperature increases, ΔH negative). Endothermic reactions absorb energy from surroundings (temperature decreases, ΔH positive). Bond breaking is endothermic; bond making is exothermic. Overall ΔH = energy to break bonds − energy released making bonds. Tags: igcse chemistry energetics Created: 2026-07-14 Last Updated: 2026-07-16

What Is Energetics?

Energetics is the study of energy changes during chemical reactions. Every chemical reaction involves energy being transferred between the reacting system and its surroundings. This energy is usually in the form of heat.

In any chemical reaction:

• Bonds in the reactants are broken — this requires energy (endothermic)
• New bonds in the products are formed — this releases energy (exothermic)

The overall energy change (ΔH, pronounced “delta H”) is the difference between these two processes.

Exothermic Reactions

An exothermic reaction is one that releases energy (usually as heat) to the surroundings.

Key characteristics:

• The temperature of the surroundings increases
• The products have less energy than the reactants
• The enthalpy change, ΔH, is negative (ΔH < 0)
• Energy is transferred from the system to the surroundings

Common examples of exothermic reactions:

Endothermic Reactions

An endothermic reaction is one that absorbs energy (usually as heat) from the surroundings.

Key characteristics:

• The temperature of the surroundings decreases
• The products have more energy than the reactants
• The enthalpy change, ΔH, is positive (ΔH > 0)
• Energy is transferred from the surroundings to the system

Common examples of endothermic reactions:

Combustion and neutralisation are the two exothermic reactions most commonly referenced. Photosynthesis and thermal decomposition are the two endothermic reactions most commonly referenced.

Energy Level Diagrams (Reaction Profiles)

Energy level diagrams show the relative energies of reactants and products during a reaction. The horizontal axis represents the progress of the reaction (also called the reaction coordinate); the vertical axis represents energy.

Exothermic Energy Level Diagram

Energy
  |    Reactants
  |       \\
  |        \\  Activation energy (Ea)
  |         \\
  |          -------
  |                 \\   ΔH (negative — energy released)
  |                  \\
  |                   \\____ Products
  |________________________________
           Progress of reaction

• Reactants start at a higher energy level than products
• The activation energy (Ea) is the minimum energy needed for the reaction to start — shown as the hump
• ΔH is negative because the products are lower in energy than the reactants
• The arrow from reactants to products points downwards

Endothermic Energy Level Diagram

Energy
  |                     /---- Products
  |                    /
  |                   /  ΔH (positive — energy absorbed)
  |                  /
  |          -------
  |         / \\
  |        /   \\  Activation energy (Ea)
  |       /     \\
  |   Reactants
  |________________________________
           Progress of reaction

• Reactants start at a lower energy level than products
• ΔH is positive because the products are higher in energy than the reactants
• The arrow from reactants to products points upwards

When drawing energy level diagrams, label: (1) the axes, (2) reactants and products, (3) Ea (activation energy), (4) ΔH with the correct sign (+ or -).

Making and Breaking Bonds

Understanding why reactions are exothermic or endothermic comes down to bond energies:

• Bond breaking is endothermic — energy must be supplied (absorbed) to overcome the attractive forces holding atoms together
• Bond making is exothermic — energy is released when new bonds form because the atoms achieve a more stable (lower energy) state

The overall energy change is the balance between these two processes:

ΔH = Σ(energy to break bonds) − Σ(energy released when making bonds)

• If more energy is released making bonds than is needed to break bonds: exothermic (ΔH negative)
• If more energy is needed to break bonds than is released making bonds: endothermic (ΔH positive)

This is covered in detail in Bond Energy Calculations.

Practical: Measuring Temperature Change

Neutralisation Experiment (HCl + NaOH)

1. Measure 25 cm3 of sodium hydroxide solution into a polystyrene cup
2. Record the initial temperature of the NaOH solution
3. Measure 25 cm3 of dilute hydrochloric acid
4. Add the acid to the cup, stir gently with the thermometer
5. Record the highest (or lowest) temperature reached
6. Calculate the temperature change

Expected result: Temperature increases by several degrees (exothermic).

Displacement Experiment (Zn + CuSO4)

1. Measure 25 cm3 of copper(II) sulfate solution into a polystyrene cup
2. Record the initial temperature
3. Add an excess of zinc powder
4. Stir and record the highest temperature reached
5. Calculate the temperature change

Expected result: Temperature increases by several degrees (exothermic).

Why a polystyrene cup? It is a good thermal insulator — it minimises heat loss to the surroundings, making the measured temperature change more accurate.

Sources of error and improvements:

• Heat loss to surroundings → use a lid on the cup
• Incomplete reaction → stir thoroughly
• Thermometer reading lag → stir and wait for steady reading

Uses of Exothermic and Endothermic Reactions

Key Points

• Exothermic: releases energy, temperature rises, ΔH negative
• Endothermic: absorbs energy, temperature falls, ΔH positive
• Exothermic examples: combustion, neutralisation, metal + acid, respiration, displacement
• Endothermic examples: photosynthesis, thermal decomposition, dissolving ammonium nitrate
• Bond breaking = endothermic (+); bond making = exothermic (−)
• ΔH = energy to break bonds − energy released forming bonds
• Energy level diagram: y-axis = energy, x-axis = progress of reaction
• Activation energy (Ea) = minimum energy needed for reaction to start
• Polystyrene cup = thermal insulator, minimises heat loss

Key Concepts from Past Papers

• Exothermic reaction: a reaction that releases heat energy to the surroundings
• Endothermic reaction: a reaction that absorbs heat energy from the surroundings
• Activation energy: the minimum energy required for a reaction to occur
• Enthalpy change (ΔH): the overall energy change in a reaction (negative for exothermic, positive for endothermic)
• Exothermic because energy level of reactants is above energy level of products
• Endothermic because heating is needed
• Bond breaking is endothermic; bond making is exothermic
• ΔH negative for exothermic; ΔH positive for endothermic

Keywords from Past Papers

exothermic, because, energy, level, endothermic, heating, needed, reactants, products, arrow, going, downwards

Related Notes

• Bond Energy Calculations
• Rates of Reaction
• IGCSE-Chem-Index

Sources

• OpenStax Chemistry 2e — [Chapter 5: Thermochemistry], Rice University (free, CC BY 4.0)
• BBC Bitesize GCSE Chemistry — [Exothermic and Endothermic Reactions], BBC (free educational resource)
• Cambridge IGCSE Chemistry 0620 — Syllabus Section 5: Chemical Energetics, Cambridge Assessment International Education
• CK-12 Chemistry for High School — [Chapter 17: Thermochemistry], CK-12 Foundation (free, CC BY-NC 3.0)

Past Paper Sources

• 0620/32 May/June 2018: Q88(d)(ii) (1m)
• 0620/33 October/November 2016: Q55(a)(i) (1m)
• 0971/32 May/June 2022: Q44(d)(i) (1m)

Common Misconceptions