GCSE — chemistryFoundation6 min read

Exothermic and Endothermic Reactions

Understand energy changes in reactions, bond energy calculations, and energy profiles for GCSE Chemistry.

Tutor AI Team2026-02-25T12:36:25.123Z

# Exothermic and Endothermic Reactions

Every chemical reaction involves energy changes. Some reactions release energy to the surroundings (you feel warmth), while others absorb energy (things get cold). Understanding these energy changes is important for GCSE Chemistry — from explaining why hand warmers work to calculating bond energies.

1. Exothermic Reactions

An exothermic reaction transfers energy to the surroundings, usually as heat. The temperature of the surroundings increases.

Key Features

Energy is released to the surroundings
Temperature of surroundings rises
$\Delta H$ is negative (energy is lost by the reaction)

Examples

Reaction	Example
Combustion	Burning fuels: $\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}$
Neutralisation	Acid + alkali: $\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$
Oxidation	Respiration, rusting
Many reactions of metals with acids	$\text{Mg} + 2\text{HCl} \rightarrow \text{MgCl}_2 + \text{H}_2$

Everyday Uses

Hand warmers (iron oxidation)
Self-heating cans (calcium oxide + water)

2. Endothermic Reactions

An endothermic reaction takes in energy from the surroundings. The temperature of the surroundings decreases.

Key Features

Energy is absorbed from the surroundings
Temperature of surroundings falls
$\Delta H$ is positive (energy is gained by the reaction)

Examples

Reaction	Example
Thermal decomposition	$\text{CaCO}_3 \rightarrow \text{CaO} + \text{CO}_2$
Photosynthesis	$6\text{CO}_2 + 6\text{H}_2\text{O} \rightarrow \text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2$
Citric acid + sodium hydrogen carbonate	Fizzy reaction that feels cold
Dissolving ammonium nitrate in water	Used in sports injury cold packs

Everyday Uses

Sports cold packs (ammonium nitrate dissolving in water)

3. Energy Profile Diagrams

Energy profile diagrams show the energy changes during a reaction.

Exothermic Profile

Reactants start at a higher energy level
Products are at a lower energy level
The difference = energy released to surroundings
There is an activation energy hill that must be overcome first

Endothermic Profile

Reactants start at a lower energy level
Products are at a higher energy level
The difference = energy absorbed from surroundings
There is still an activation energy to overcome

Activation Energy

The activation energy ( $E_a$ ) is the minimum energy that particles must have to react when they collide. It's shown as the "hill" on the energy profile diagram.

All reactions require activation energy
A catalyst lowers the activation energy (provides an alternative reaction pathway)
A match provides activation energy to start combustion

4. Bond Energy Calculations

The Principle

Breaking bonds requires energy → endothermic (energy IN)
Making bonds releases energy → exothermic (energy OUT)

The overall energy change depends on the balance:

$\Delta H = \text{energy to break bonds} - \text{energy released making bonds}$

If more energy is released making bonds than needed to break them → exothermic ( $\Delta H$ negative)
If more energy is needed to break bonds than released making them → endothermic ( $\Delta H$ positive)

Bond Energy Values

Bond energies tell you how much energy is needed to break one mole of a particular bond.

Bond	Bond Energy (kJ/mol)
C−H	413
C−C	347
C=C	614
O=O	498
O−H	464
C=O	805
H−H	436
C−O	358
N≡N	945
H−Cl	432

Worked Example: Bond Energy Calculation

Problem

Question: Calculate the energy change for the combustion of methane: $\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}$

Bonds broken (reactants):

4 × C−H = $4 \times 413 = 1652$ kJ
2 × O=O = $2 \times 498 = 996$ kJ
Total energy in = $1652 + 996 = 2648$ kJ

Bonds made (products):

2 × C=O = $2 \times 805 = 1610$ kJ
4 × O−H = $4 \times 464 = 1856$ kJ
Total energy out = $1610 + 1856 = 3466$ kJ

Overall: $\Delta H = 2648 - 3466 = -818 \text{ kJ/mol}$

Negative → exothermic ✓ (combustion releases energy)

Solution

Worked Example: Interpreting a Profile Diagram

Problem

Question: A reaction has reactants at 200 kJ and products at 350 kJ. Is it exothermic or endothermic? What is the energy change?

Solution

Products are at a higher energy than reactants → endothermic
Energy change = $350 - 200 = +150$ kJ

Worked Example: Required Practical

Problem

Question: A student added magnesium to hydrochloric acid. The temperature rose from 21°C to 35°C. Is this exothermic or endothermic?

Solution

The temperature of the surroundings (the solution) increased, so the reaction transferred energy to the surroundings. This is exothermic.

6. Required Practical: Investigating Temperature Changes

Method

Measure 25 cm³ of acid into a polystyrene cup (insulation)
Record the initial temperature
Add a measured amount of solid or solution
Stir and record the temperature every 30 seconds
Find the maximum (or minimum) temperature
Calculate the temperature change: $\Delta T = T_{\text{final}} - T_{\text{initial}}$

Why Polystyrene Cup?

Polystyrene is a good insulator — it minimises heat loss to the environment, giving more accurate results.

7. Practice Questions

1. Define exothermic and endothermic reactions.
1. Classify these as exothermic or endothermic: (a) combustion, (b) photosynthesis, (c) neutralisation, (d) thermal decomposition.
1. Draw energy profile diagrams for an exothermic and endothermic reaction, labelling reactants, products, activation energy, and overall energy change.
1. Calculate the energy change for: $\text{H}_2 + \text{Cl}_2 \rightarrow 2\text{HCl}$ Given: H−H = 436 kJ/mol, Cl−Cl = 243 kJ/mol, H−Cl = 432 kJ/mol
1. Explain what activation energy is and how a catalyst affects it.

Want to check your answers and get step-by-step solutions?

8. Common Misconceptions

Misconception	Reality
Exothermic reactions don't need energy to start	All reactions need activation energy to start
Making bonds requires energy	Making bonds releases energy; breaking bonds requires energy
Temperature drop = no reaction	An endothermic reaction is still a reaction — it just absorbs energy
Energy is created in exothermic reactions	Energy is transferred from chemical stores, not created (conservation of energy)

9. Exam Tips

"Exo" = exit = energy goes OUT; "Endo" = enter = energy comes IN
Bond energy calculations: BREAK bonds first (energy in), then MAKE bonds (energy out)
If $\Delta H$ is negative = exothermic; positive = endothermic
On energy profiles, the activation energy is measured from reactants to the peak
For the practical, always mention insulation (polystyrene cup) and why

Summary

Exothermic: energy released, temperature rises, $\Delta H < 0$
Endothermic: energy absorbed, temperature falls, $\Delta H > 0$
Breaking bonds = endothermic; Making bonds = exothermic
$\Delta H = \text{energy in (breaking)} - \text{energy out (making)}$
Activation energy = minimum energy needed for a reaction to start
Catalysts lower the activation energy

Tags:

#gcse#chemistry#energy-changes#exothermic#endothermic

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# Exothermic and Endothermic Reactions

1. Exothermic Reactions

Key Features

Examples

Everyday Uses

2. Endothermic Reactions

Key Features

Examples

Everyday Uses

3. Energy Profile Diagrams

Exothermic Profile

Endothermic Profile

Activation Energy

4. Bond Energy Calculations

The Principle

Bond Energy Values

Worked Example: Bond Energy Calculation

Worked Example: Interpreting a Profile Diagram

Worked Example: Required Practical

6. Required Practical: Investigating Temperature Changes

Method

Why Polystyrene Cup?

7. Practice Questions

8. Common Misconceptions

9. Exam Tips

Summary

Related Topics

Evolution and Natural Selection

Biodiversity and Conservation

Sampling and Field Investigations

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