Exothermic and Endothermic Reactions

Energy changes are a fundamental part of every chemical reaction. In this lesson you will learn the difference between exothermic and endothermic reactions, understand how energy is transferred to and from the surroundings, and explore everyday examples of both types. This is a core topic within the Energy Changes chapter of the AQA GCSE Combined Science Trilogy specification (8464).

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What Happens to Energy During a Chemical Reaction?

During any chemical reaction, bonds in the reactants are broken and new bonds are formed in the products. Both processes involve energy changes:

Process	Energy Change
Breaking bonds	Energy is taken in from the surroundings (endothermic)
Making bonds	Energy is released to the surroundings (exothermic)

The overall energy change of a reaction depends on the balance between these two processes:

$\Delta H = \text{Energy to break bonds} - \text{Energy released making bonds}$ΔH=Energy to break bonds−Energy released making bonds

Exam Tip: A very common mistake is to say that breaking bonds releases energy. Remember: breaking bonds ALWAYS requires energy. Only when NEW bonds form is energy released.

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Exothermic Reactions

An exothermic reaction is one that transfers energy to the surroundings. The temperature of the surroundings increases.

In an exothermic reaction:

$\text{Energy released (making bonds)} > \text{Energy taken in (breaking bonds)}$Energy released (making bonds)>Energy taken in (breaking bonds)

Therefore $\Delta H$ΔH is negative.

Examples of Exothermic Reactions

Reaction Type	Example
Combustion	Burning methane: $\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}$CH4​+2O2​→CO2​+2H2​O
Neutralisation	$\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$HCl+NaOH→NaCl+H2​O
Oxidation	Rusting of iron, respiration
Displacement	Magnesium reacting with copper sulfate solution

Everyday Uses of Exothermic Reactions

• Self-heating cans — calcium oxide reacts with water in a compartment, heating the contents.
• Hand warmers — iron powder is slowly oxidised, releasing heat over several hours.
• Combustion in engines — burning petrol or diesel powers vehicles.

graph LR
    A[Reactants] -->|"Bonds broken \n(energy taken in)"| B[Transition State]
    B -->|"New bonds formed \n(energy released)"| C[Products]
    C -->|"MORE energy released \nthan taken in"| D["Surroundings get HOTTER \n(ΔH is negative)"]
    style D fill:#ff9999,stroke:#cc0000

Exam Tip: When a question asks you to identify an exothermic reaction, look for clues such as a temperature rise, flames, or the word "combustion." Neutralisation reactions are also always exothermic — this is a favourite AQA exam question.

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Endothermic Reactions

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

In an endothermic reaction:

$\text{Energy taken in (breaking bonds)} > \text{Energy released (making bonds)}$Energy taken in (breaking bonds)>Energy released (making bonds)

Therefore $\Delta H$ΔH is positive.

Examples of Endothermic Reactions

Reaction Type	Example
Thermal decomposition	Heating calcium carbonate: $\text{CaCO}_3 \rightarrow \text{CaO} + \text{CO}_2$CaCO3​→CaO+CO2​
Citric acid + sodium hydrogencarbonate	Classic endothermic reaction in solution
Photosynthesis	$6\text{CO}_2 + 6\text{H}_2\text{O} \xrightarrow{\text{light}} \text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2$6CO2​+6H2​Olight​C6​H12​O6​+6O2​
Dissolving ammonium nitrate in water	Solution becomes noticeably cold

Everyday Uses of Endothermic Reactions

• Instant cold packs — ammonium nitrate dissolves in water inside the pack, absorbing heat and cooling the injured area.
• Baking — thermal decomposition of sodium hydrogencarbonate (baking soda) helps dough rise.

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Comparing Exothermic and Endothermic Reactions

Feature	Exothermic	Endothermic
Energy transfer	To the surroundings	From the surroundings
Temperature change	Surroundings get hotter	Surroundings get cooler
Sign of $\Delta H$ΔH	Negative	Positive
Bond energy balance	Energy out > energy in	Energy in > energy out
Examples	Combustion, neutralisation, oxidation	Thermal decomposition, photosynthesis
Everyday use	Hand warmers, self-heating cans	Cold packs

graph TD
    subgraph "Exothermic"
        A1["Energy released > Energy absorbed"] --> B1["ΔH is NEGATIVE"]
        B1 --> C1["Temperature RISES"]
    end
    subgraph "Endothermic"
        A2["Energy absorbed > Energy released"] --> B2["ΔH is POSITIVE"]
        B2 --> C2["Temperature FALLS"]
    end

Exam Tip: The prefixes help you remember. "Exo" means out (energy goes OUT to surroundings). "Endo" means in (energy comes IN from surroundings). Write these prefixes in your revision notes.

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Reversible Reactions and Energy Changes

Some reactions are reversible, meaning they can go in both directions. When a reversible reaction is exothermic in one direction, it is endothermic in the reverse direction, and the energy transferred is the same magnitude in both cases.

Example — hydrated and anhydrous copper sulfate:

$\text{CuSO}_4 \cdot 5\text{H}_2\text{O} \xrightleftharpoons[\text{add water (exothermic)}]{\text{heat (endothermic)}} \text{CuSO}_4 + 5\text{H}_2\text{O}$CuSO4​⋅5H2​Oheat (endothermic)add water (exothermic)​CuSO4​+5H2​O

graph LR
    A["Hydrated copper sulfate \n(blue crystals)"] -->|"Heat — endothermic"| B["Anhydrous copper sulfate \n(white powder) + Water"]
    B -->|"Add water — exothermic"| A

• Heating hydrated copper sulfate (blue) produces anhydrous copper sulfate (white) — endothermic.
• Adding water to anhydrous copper sulfate turns it blue again — exothermic.

Exam Tip: AQA (8464) often asks about reversible reactions and energy. Remember: the energy change in the forward direction is equal and opposite to the energy change in the reverse direction.

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Common Mistakes

Mistake	Correction
"Breaking bonds releases energy"	Breaking bonds always REQUIRES energy (endothermic)
Confusing exothermic and endothermic	Exo = out, Endo = in — link to temperature change
Forgetting reversible reaction energy rule	Forward exothermic = reverse endothermic, same magnitude
Thinking all reactions need heat to start	Exothermic reactions still need activation energy to begin

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Summary

• Exothermic reactions transfer energy to the surroundings, causing a temperature rise (e.g., combustion, neutralisation). $\Delta H$ΔH is negative.
• Endothermic reactions take in energy from the surroundings, causing a temperature drop (e.g., thermal decomposition, dissolving ammonium nitrate). $\Delta H$ΔH is positive.
• Breaking bonds is endothermic; forming bonds is exothermic.
• The overall energy change depends on whether more energy is released forming bonds or taken in breaking bonds.
• In reversible reactions, the energy change in the forward direction is equal and opposite to that in the reverse direction.
• Energy is always conserved in chemical reactions.

Exam Tip (AQA 8464): A 6-mark question on exothermic vs endothermic reactions is very common. Make sure you can clearly define both terms, give examples of each, explain the role of bond breaking and forming, and link temperature changes to the type of reaction.

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Deeper Dive: Why Breaking Bonds Is Endothermic

Covalent bonds exist because two atoms share electrons to reach a lower, more stable energy state than when separated. To pull those atoms apart, we have to put energy in to overcome the electrostatic attraction between the shared electrons and the positive nuclei. The amount of energy needed for one mole of a specific bond is called the bond dissociation energy (also known as the bond energy), and it is always a positive quantity.

When new bonds form, the reverse is true: atoms fall into a lower energy state and the excess energy is released to the surroundings — usually as heat, but sometimes as light, sound, or electrical energy. This is why forming bonds is always exothermic.

Step	Process	Energy Direction	Sign
1	Reactant bonds break	Energy absorbed from surroundings	+ (endothermic step)
2	Product bonds form	Energy released to surroundings	− (exothermic step)
Overall	Net change = step 1 + step 2	Depends on magnitudes	+ or −

Common mistake callout: Students often write "the reactants release energy when they break apart." This is wrong and will lose marks. Always say "energy is taken in to break the reactant bonds, and energy is released when the new product bonds form."

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Worked Example: Predicting Sign of $\Delta H$ΔH from Bond Energies

Question: For the reaction $\text{H}_2 + \text{Cl}_2 \rightarrow 2\text{HCl}$H2​+Cl2​→2HCl, the bond energies are H–H = 436 kJ/mol, Cl–Cl = 242 kJ/mol, H–Cl = 432 kJ/mol. Is this reaction exothermic or endothermic?

Step 1 — Energy to break bonds (reactants):

$E_{\text{in}} = 436 + 242 = 678 \text{ kJ/mol}$Ein​=436+242=678 kJ/mol

Step 2 — Energy released forming bonds (products):

$E_{\text{out}} = 2 \times 432 = 864 \text{ kJ/mol}$Eout​=2×432=864 kJ/mol

Step 3 — Calculate $\Delta H$ΔH:

$\Delta H = E_{\text{in}} - E_{\text{out}} = 678 - 864 = -186 \text{ kJ/mol}$ΔH=Ein​−Eout​=678−864=−186 kJ/mol

Step 4 — Interpret the sign:

The value is negative, so the reaction is exothermic. More energy is released when the two H–Cl bonds form than is needed to break the H–H and Cl–Cl bonds. The surroundings would get warmer.

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Spotting Exothermic and Endothermic Reactions in Unfamiliar Contexts

AQA often gives you an unfamiliar reaction and asks you to identify whether it is exothermic or endothermic. Use these clues:

Clue in the Question	Likely Answer
"Temperature of the mixture rose from 20 °C to 32 °C"	Exothermic
"The beaker felt cold to the touch"	Endothermic
"Light and heat were given out"	Exothermic (combustion)
"Energy had to be continuously supplied for the reaction to continue"	Endothermic
"$\Delta H$ΔH = −245 kJ/mol"	Exothermic (negative sign)
"The reaction mixture drew heat from the surroundings"	Endothermic
"The reaction is the reverse of an exothermic reaction"	Endothermic (same magnitude, opposite sign)

graph TD
    A["Look at the clue in the question"] --> B{"Does the temperature RISE or FALL, or is ΔH given?"}
    B -->|"Temperature rises / ΔH negative"| C["EXOTHERMIC — energy released"]
    B -->|"Temperature falls / ΔH positive"| D["ENDOTHERMIC — energy absorbed"]
    C --> E["Examples: combustion, neutralisation, oxidation, displacement"]
    D --> F["Examples: thermal decomposition, dissolving ammonium nitrate, photosynthesis"]

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Neutralisation: Why It Is Always Exothermic

Neutralisation reactions between an acid and an alkali are always exothermic. The underlying ionic equation is:

$\text{H}^+ (aq) + \text{OH}^- (aq) \rightarrow \text{H}_2\text{O} (l)$H+(aq)+OH−(aq)→H2​O(l)

Strong O–H bonds form in water, releasing about 57 kJ of energy per mole of water made. Because this ionic equation is the same for any strong acid reacting with a strong alkali, the $\Delta H$ΔH of neutralisation is approximately −57 kJ/mol regardless of which strong acid and strong alkali are used.

Exam Tip (AQA 8464): If you are told "0.1 mol of H⁺ reacted with 0.1 mol of OH⁻" and asked to estimate the heat released, multiply: $0.1 \times 57 = 5.7$0.1×57=5.7 kJ. This is a common Higher Tier calculation.

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Answering at different grade levels

A student aiming for grades 4–5 should be able to:

• State that an exothermic reaction transfers energy to the surroundings and that the surroundings get hotter.
• State that an endothermic reaction takes energy in from the surroundings and that the surroundings get cooler.
• Give one or two everyday examples of each (combustion, neutralisation for exothermic; thermal decomposition, dissolving ammonium nitrate for endothermic).
• Identify the reaction type from a simple temperature observation.

A student aiming for grades 6–7 should additionally:

• Explain that breaking bonds is endothermic (requires energy input) and forming bonds is exothermic (releases energy).
• Describe that the overall energy change depends on the balance between these two processes.
• State that in a reversible reaction the energy change in the forward direction is equal and opposite to that in the reverse direction.
• Correctly use the terms activation energy and reaction profile when describing examples.

A student aiming for grades 8–9 should additionally:

• Predict the sign of $\Delta H$ΔH from bond energies by calculating $E_{\text{bonds broken}} - E_{\text{bonds made}}$Ebonds broken​−Ebonds made​.
• Explain why neutralisation between any strong acid and strong alkali releases approximately the same energy per mole.
• Apply the concepts of exothermic and endothermic reactions to unfamiliar contexts, linking temperature observations quantitatively to $\Delta H$ΔH and to the underlying bond changes.
• Use precise scientific vocabulary (exothermic, endothermic, activation energy, bond energy, reaction profile) throughout an extended written answer.

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AQA alignment: This content is aligned with AQA GCSE Combined Science: Trilogy (8464) specification section 5.5 Energy changes — specifically 5.5.1.1 Exothermic and endothermic reactions. Assessed on Chemistry Paper 1.