What are energy changes in chemical reactions?

Energy changes in chemical reactions refer to the energy absorbed or released during a chemical process. In the Cambridge IGCSE Coordinated Science curriculum, this concept is crucial as it helps explain how bonds are broken and formed, leading to either an exothermic or endothermic reaction.

How do exothermic and endothermic reactions differ in terms of energy changes?

Exothermic reactions release energy to the surroundings, usually in the form of heat, resulting in an increase in temperature. In contrast, endothermic reactions absorb energy from the surroundings, leading to a decrease in temperature. Understanding these differences is essential for Cambridge IGCSE students studying energy changes in chemical reactions.

Why is it important to study energy changes in chemical reactions in Coordinated Science?

Studying energy changes in chemical reactions is important because it helps students understand how energy is conserved and transformed during chemical processes. This knowledge is fundamental in predicting reaction behavior and is a key component of the Cambridge IGCSE Coordinated Science syllabus.

What role do bond energies play in energy changes during chemical reactions?

Bond energies are crucial in determining the energy changes during chemical reactions. Breaking bonds requires energy, while forming bonds releases energy. The difference between these energies determines whether a reaction is exothermic or endothermic. This concept is a vital part of the Cambridge IGCSE Coordinated Science curriculum.

Can you provide an example of an exothermic and an endothermic reaction?

An example of an exothermic reaction is the combustion of methane, where energy is released as heat. An example of an endothermic reaction is the thermal decomposition of calcium carbonate, where energy is absorbed. These examples help illustrate the concept of energy changes in chemical reactions, as taught in the Cambridge IGCSE Coordinated Science course.

Why is "Saying bond breaking is exothermic" a common mistake in Energy Changes in Chemical Reactions?

Why it happens: Students confuse the direction of energy for bond breaking and bond forming. How to avoid it: Bond BREAKING = endothermic (energy IN required). Bond FORMING = exothermic (energy OUT released). Use the mnemonic: 'Breaking Bad = absorbs energy; Forming Friends = releases energy.' Source: 0654 Examiner Report 2022

Why is "Saying a catalyst changes the overall ΔH of a reaction" a common mistake in Energy Changes in Chemical Reactions?

Why it happens: Students think 'easier pathway' means 'different energy products or reactants'. How to avoid it: A catalyst lowers the activation energy but does NOT change the energy levels of reactants or products; ΔH is the same with or without a catalyst.

Why is "Counting CO₂ as having one C=O bond rather than two" a common mistake in Energy Changes in Chemical Reactions?

Why it happens: Students see one carbon dioxide formula and assume one bond. How to avoid it: CO₂ is O=C=O with TWO C=O double bonds. Multiply 805 kJ/mol by 2 for each CO₂ molecule, and by 4 total if there are 2 CO₂ molecules. Source: 0654 Examiner Report 2023

Energy Changes in Chemical ReactionsCambridge IGCSE Coordinated Science (0654)

Energy Changes in Chemical Reactions

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Energy Changes in Chemical Reactions — Cambridge IGCSE Coordinated Science 0654

Chemical reactions involve energy changes. Exothermic reactions release energy to surroundings (temperature rises). Endothermic reactions absorb energy (temperature falls). Cambridge tests energy profile diagrams and activation energy.

What you’ll learn

Mapped to the Cambridge IGCSE 0654 syllabus (2025-2027).

C6.1 — Define exothermic and endothermic reactions.
C6.2 — Draw and interpret energy profile (reaction pathway) diagrams.
C6.3 — Define activation energy and explain the effect of catalysts on Ea.
C6.4 — Calculate energy changes from bond energies.

Exothermic and Endothermic Reactions

Exothermic reactions release energy (temperature rises). Endothermic reactions absorb energy (temperature falls).

Exothermic reaction: a reaction that releases energy (usually as heat) to the surroundings. Temperature of surroundings RISES.

ΔH is NEGATIVE (products have less energy than reactants)
Bond forming releases more energy than bond breaking absorbs

Endothermic reaction: a reaction that absorbs energy from the surroundings. Temperature of surroundings FALLS.

ΔH is POSITIVE (products have more energy than reactants)
Bond breaking absorbs more energy than bond forming releases

Common examples:

Exothermic	Endothermic
Combustion of fuels	Thermal decomposition of CaCO₃
Neutralisation (acid + base)	Photosynthesis
Respiration	Dissolving ammonium nitrate
Rusting (oxidation of iron)	Melting, boiling (physical endothermic)
Hand warmers	Cold packs

Measuring energy changes: Bomb calorimeter or simple calorimetry (temperature change in water). Energy released (J) = mass of water × specific heat capacity (4.18 J/g°C) × ΔT

Exothermic: T rises, ΔH negative (combustion, neutralisation, respiration).
Endothermic: T falls, ΔH positive (thermal decomposition, photosynthesis, dissolving NH₄NO₃).
Energy = m × c × ΔT (calorimetry calculation).

Energy Profile Diagrams

Energy profile diagrams show the energy of reactants, products, transition state, and activation energy — for both exothermic and endothermic reactions.

Energy profile diagram (reaction pathway diagram): a graph showing how the energy of the system changes as reactants become products.

Key features:

Y-axis: energy; X-axis: reaction progress (coordinate)
Reactants: starting energy level
Products: ending energy level
Transition state (activated complex): peak of curve — highest energy point
Activation energy (Ea): energy difference between reactants and the peak (top of hump)
ΔH (enthalpy change): energy difference between reactants and products

Exothermic profile: products are at LOWER energy than reactants. ΔH negative. Hump still present (Ea).

Endothermic profile: products are at HIGHER energy than reactants. ΔH positive. Hump still present.

Effect of catalysts on the energy profile:

Catalyst provides an ALTERNATIVE reaction pathway with LOWER activation energy
Shows as a LOWER hump on the diagram
ΔH is unchanged (reactant and product energy levels stay the same)
More reactant molecules have enough energy to react → faster reaction

Exothermic: products lower than reactants (ΔH negative). Endothermic: products higher (ΔH positive).
Activation energy = energy from reactants to peak of curve.
Catalyst: lowers Ea (lower hump) but does NOT change ΔH.

Bond Energies — Calculating Energy Changes

Energy is absorbed when bonds break; released when bonds form. ΔH = energy in (breaking) − energy out (forming).

Bond energy: the energy required to break one mole of a specified bond in gaseous molecules. Units: kJ/mol.

Principle:

Breaking bonds: ENDOTHERMIC (energy ABSORBED)
Forming bonds: EXOTHERMIC (energy RELEASED)

Calculating ΔH from bond energies:

ΔH = energy absorbed (bonds broken) − energy released (bonds formed)

If ΔH is negative → exothermic (more energy released forming bonds than absorbed breaking them)
If ΔH is positive → endothermic (more energy absorbed breaking bonds than released forming them)

Example — Combustion of hydrogen: H₂ + ½O₂ → H₂O

Bonds broken: 1 × H–H (436 kJ) + ½ × O=O (248 kJ) = 684 kJ
Bonds formed: 2 × O–H (2 × 463 = 926 kJ)
ΔH = 684 − 926 = −242 kJ/mol (exothermic)

Average bond energies: values given in exam data sheets (or must be recalled for common bonds):

C–H: ~413, C–C: ~347, C=C: ~612, O=O: ~498, O–H: ~463, H–H: ~436

Breaking bonds: absorbs energy (endothermic). Forming bonds: releases energy (exothermic).
ΔH = total energy IN (bonds broken) − total energy OUT (bonds formed).
Negative ΔH = exothermic. Positive ΔH = endothermic.

How it’s examined

Paper 4: Draw an energy profile diagram for an exothermic reaction, labelling reactants, products, Ea, and ΔH (4 marks). 'Use the bond energies to calculate the enthalpy change for this reaction' (3 marks — must show bonds broken, bonds formed, correct arithmetic). 'Explain how a catalyst affects the activation energy, and hence the reaction rate' (3 marks — lower Ea, more molecules have sufficient energy, faster rate). MCQ tests exo/endothermic identification from described reactions.

Sources: Cambridge IGCSE Coordinated Sciences 0654 syllabus 2025-2027 (C6); 0654 Examiner Reports 2022-2024. Last reviewed 2026-05-14.

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Step-by-step worked examples — Energy Changes in Chemical Reactions

Step-by-step solutions to past-paper-style questions on energy changes in chemical reactions, written exactly the way a tutor would explain them at the board.

1Bond energy calculation for combustion of methane
Extended
Question
The combustion of methane: $\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}$ . Bond energies: C–H = 413; O=O = 498; C=O = 805; O–H = 464 kJ/mol. Calculate $\Delta H$ for this reaction.
Step-by-step solution
1. Step 1
  Identify bonds broken (reactants).
  $\text{Bonds broken} = 4(\text{C--H}) + 2(\text{O=O}) = 4(413) + 2(498) = 1652 + 996 = 2648\,\text{kJ/mol}$
2. Step 2
  Identify bonds formed (products).
  $\text{Bonds formed} = 2(\text{C=O in CO}_2) + 4(\text{O--H in H}_2\text{O}) = 2(2 \times 805) + 4(464) = 3220 + 1856\,\text{kJ/mol}$
3. Step 3
  Note: each CO₂ has 2 C=O bonds, so 2 CO₂ molecules = 4 C=O bonds total.
  $\text{Bonds formed} = 4(805) + 4(464) = 3220 + 1856 = 5076\,\text{kJ/mol}$
4. Step 4
  Calculate ΔH.
  $\Delta H = \text{energy in} - \text{energy out} = 2648 - 5076 = -2428\,\text{kJ/mol}$
Answer
$\Delta H = -2428\,\text{kJ/mol}$ (exothermic)
Examiner tip
The negative sign confirms exothermic. Remember: CO₂ has TWO double bonds (O=C=O); do not count it as one bond.
2Describing an exothermic reaction
Core
Question
Describe what is meant by an exothermic reaction and give TWO examples.
Step-by-step solution
1. Step 1
  An exothermic reaction releases energy (heat) to the surroundings; the temperature of the surroundings increases; ΔH is negative.
2. Step 2
  Example 1: combustion of fuels (e.g. burning methane — produces heat and light).
3. Step 3
  Example 2: neutralisation of an acid with an alkali (temperature of solution rises).
Answer
Exothermic: releases heat to surroundings; temperature rises; ΔH < 0. Examples: combustion, neutralisation.
Examiner tip
Other valid examples include: respiration, rusting, oxidation reactions, dissolving in concentrated acids.
3Effect of catalyst on activation energy
Extended
Question
Explain, using an energy profile diagram, how a catalyst speeds up a chemical reaction.
Step-by-step solution
1. Step 1
  In an energy profile diagram, the activation energy is the energy barrier between reactants and products.
2. Step 2
  A catalyst provides an alternative reaction pathway with a lower activation energy.
3. Step 3
  With a lower activation energy, a greater proportion of colliding particles have sufficient energy to react (exceed the new, lower activation energy).
4. Step 4
  More successful collisions per second → faster rate. The overall energy change (ΔH) is unchanged; the catalyst is not consumed.
Answer
Catalyst lowers activation energy → more particles have energy ≥ Ea → more successful collisions per unit time → faster reaction; ΔH unchanged; catalyst not consumed.
Examiner tip
Examiners require: 'lower activation energy', 'more particles have sufficient energy' (or 'more successful collisions'). Simply saying 'speeds up reaction' without explaining why earns no marks.
4Dissolving ammonium nitrate — endothermic evidence
Core
Question
State whether dissolving ammonium nitrate in water is exothermic or endothermic and describe the evidence.
Step-by-step solution
1. Step 1
  Dissolving ammonium nitrate in water is endothermic.
2. Step 2
  Evidence: the temperature of the solution decreases; the surroundings feel cold; a thermometer shows a drop in temperature.
Answer
Endothermic — the temperature of the water decreases as ammonium nitrate dissolves, showing that energy is absorbed from the surroundings.

Key Formulae — Energy Changes in Chemical Reactions

The formulae you need to memorise for energy changes in chemical reactions on the Cambridge IGCSE paper, with every variable defined in plain English and a note on when to use it.

Bond energy equation
$\Delta H = \sum E(\text{bonds broken}) - \sum E(\text{bonds formed})$
$\sum E(\text{bonds broken})$
total energy required to break all bonds in reactants (endothermic, +ve)
$\sum E(\text{bonds formed})$
total energy released when bonds form in products (exothermic, −ve contribution)
When to use
Calculating the enthalpy change of a reaction from bond energies. A negative ΔH = exothermic; positive ΔH = endothermic.

Key Definitions and Keywords — Energy Changes in Chemical Reactions

Definitions to memorise and the exact keywords mark schemes credit for energy changes in chemical reactions answers — sharpened from recent examiner reports for the 2026 Cambridge IGCSE sitting.

Exothermic reaction
Examiner keyword
A reaction that releases energy (heat) to the surroundings; the temperature of the surroundings increases; ΔH is negative.
Endothermic reaction
Examiner keyword
A reaction that absorbs energy (heat) from the surroundings; the temperature of the surroundings decreases; ΔH is positive.
Activation energy
Examiner keyword
The minimum energy that colliding particles must possess for a reaction to occur; shown as the peak on an energy profile diagram.
Catalyst
Examiner keyword
A substance that increases the rate of a chemical reaction by providing an alternative pathway with a lower activation energy; it is not consumed in the reaction.
Enthalpy change (ΔH)
Examiner keyword
The heat energy transferred to or from the surroundings during a chemical reaction at constant pressure. Units: kJ/mol. Negative for exothermic; positive for endothermic.
Bond energy
The energy required to break one mole of a specified covalent bond in a gaseous molecule; also equal to the energy released when that bond forms.

Common Mistakes and Misconceptions — Energy Changes in Chemical Reactions

The traps other students keep falling into on energy changes in chemical reactions questions — taken from recent Cambridge IGCSE examiner reports and mark schemes — and how to avoid them.

✕Saying bond breaking is exothermic
0654 Examiner Report 2022
Why it happens
Students confuse the direction of energy for bond breaking and bond forming.
How to avoid it
Bond BREAKING = endothermic (energy IN required). Bond FORMING = exothermic (energy OUT released). Use the mnemonic: 'Breaking Bad = absorbs energy; Forming Friends = releases energy.'
✕Saying a catalyst changes the overall ΔH of a reaction
Why it happens
Students think 'easier pathway' means 'different energy products or reactants'.
How to avoid it
A catalyst lowers the activation energy but does NOT change the energy levels of reactants or products; ΔH is the same with or without a catalyst.
✕Counting CO₂ as having one C=O bond rather than two
0654 Examiner Report 2023
Why it happens
Students see one carbon dioxide formula and assume one bond.
How to avoid it
CO₂ is O=C=O with TWO C=O double bonds. Multiply 805 kJ/mol by 2 for each CO₂ molecule, and by 4 total if there are 2 CO₂ molecules.

Practice questions

Exam-style questions with step-by-step worked solutions. Try one before checking the method.

Past paper style quiz

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Energy Changes in Chemical Reactions — frequently asked questions

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Energy Changes in Chemical Reactions

Short Study Notes in the form of Slides

Detailed Study Notes

What you’ll learn

How it’s examined

1Bond energy calculation for combustion of methane

2Describing an exothermic reaction

3Effect of catalyst on activation energy

4Dissolving ammonium nitrate — endothermic evidence

Bond energy equation

Exothermic reaction

Endothermic reaction

Activation energy

Catalyst

Enthalpy change (ΔH)

Bond energy

✕Saying bond breaking is exothermic

✕Saying a catalyst changes the overall ΔH of a reaction

✕Counting CO₂ as having one C=O bond rather than two

Practice questions

Past paper style quiz

Assess your understanding

Energy Changes in Chemical Reactions — frequently asked questions