✓ Revise Right course
Energy change = energy needed to break bonds in the reactants - energy released forming bonds in the products. Free notes and 9 questions answered.
Calculating Energy Changes Using Bond Energies is part of Exothermic and endothermic reactions in GCSE AQA Higher Chemistry.
4 min
Reading time
11
Lessons in this topic
GCSE AQA Higher
Pathway
Reviewed by Revise Right
Content review
Energy change = energy needed to break bonds in the reactants - energy released forming bonds in the products. Free notes and 9 questions answered.
AQA GCSE Chemistry mock draft, 26–31 Oct. Draft coverage: Chemistry Paper 1 (Higher tier). Reviewed for an original-practice sitting; check student availability before registering. Marking and results will follow later, after separate approval. Eligible new accounts receive a five-day Pro trial when signup email verification completes. Previous trials keep their original expiry. After expiry, new papers need Pro unless you have an explicitly confirmed free sitting offer for this event. That offer adds no Pro time or AI learning. Saved attempts and released results remain readable. No card or automatic charge.
Use this lesson to stay inside Exothermic and endothermic reactions while connecting the detail back to the wider Topic 5: Energy changes topic.
Read the core explanations from Calculating Energy Changes Using Bond Energies before testing yourself from memory.
Every chemical reaction involves breaking old bonds and making new ones. But here's the thing - breaking bonds always requires energy, whilst forming bonds always releases energy. The balance between these two processes determines whether a reaction heats up your surroundings or cools them down.
Think of it like demolishing an old house and building a new one. You need energy to knock down the old structure, but you get some energy back from selling the materials to build something new. Whether you end up richer or poorer depends on the difference between what you spent and what you gained.
Every reaction follows the same pattern. First, you break the bonds in your reactants - this always requires energy input. Then, you form new bonds in your products - this always releases energy. The overall energy change is the difference between these two steps.
If Energy In > Energy Out, the reaction is endothermic (absorbs energy from surroundings). If Energy Out > Energy In, the reaction is exothermic (releases energy to surroundings). The bigger the difference, the more dramatic the temperature change.
Let's tackle a classic exam question. The method is always the same: identify bonds broken, identify bonds formed, look up the energies, then subtract. This catches many students out because they forget to count how many of each bond type there are.
Worked example: Calculating Energy Change for Methane Combustion
Calculate the energy change for: CH₄ + 2O₂ → CO₂ + 2H₂O Bond energies: C-H = 413 kJ/mol, O=O = 498 kJ/mol, C=O = 805 kJ/mol, O-H = 464 kJ/mol
Answer: -818 kJ/mol (exothermic reaction)
Notice the negative sign tells us this reaction releases energy. That's why methane burns so well as a fuel - it gives out lots of energy to heat your home. The bigger the negative number, the more energy released.
Students lose marks on these questions for predictable reasons. The most common error is miscounting bonds - especially in molecules like water where you might forget there are two O-H bonds per molecule, not one.
Another trap is the sign convention. Bond energy tables give you positive values, but you need to think about whether you're breaking bonds (energy in) or forming bonds (energy out). The formula Energy In - Energy Out handles the signs for you.
Bond energy questions follow a predictable structure in exams. You'll get the balanced equation, a table of bond energies, and need to calculate the overall energy change. The mark scheme rewards clear method as much as the right answer.
For the full 4-6 marks available, you must show: identification of bonds broken and formed, correct use of bond energy values, proper calculation of energy in and energy out, and the final subtraction with correct sign and units.
Exam precision note: In Calculating Energy Changes Using Bond Energies - The energy change of reactions (HT only) - Exothermic and endothermic reactions, always identify the key chemical idea first, then link it to evidence, data, or an equation. A high-scoring GCSE answer names the particles involved, states what changes, and justifies the conclusion using correct scientific language. Where calculations are used, show each step with units and keep significant figures sensible. Finally, check that terminology matches AQA wording so your answer is unambiguous and mark-scheme friendly.
Bond-energy calculations compare the energy needed to break bonds with the energy released when new bonds form. If more energy is released than taken in, the reaction is exothermic. If more energy is needed to break bonds than is released when new bonds form, the reaction is endothermic.
Published by Revise Right. Use the official specification reference to check the requirements for your course. Content updated: 2026-03-19.
Move to the related lessons or back to the topic page when you need the wider sequence before exam-style practice.
Energy change = energy needed to break bonds in the reactants (energy in) - energy released forming bonds in the products (energy out). Identify every bond broken and formed, count how many of each there are, add up their bond energies, then subtract.
Bond energies are measured in kJ/mol, because each value is the energy needed to break one mole of a particular covalent bond in the gaseous state. The overall energy change of a reaction is also given in kJ/mol.
A negative energy change means the reaction is exothermic, because more energy is released forming new bonds than is needed to break the old bonds. For example, methane combustion has an energy change of -818 kJ/mol; the bigger the negative number, the more energy released.
The most common error in bond energy calculations is miscounting bonds, such as counting one O-H bond per water molecule instead of two. Another trap is the sign: the formula energy in - energy out handles the signs, giving a negative value for exothermic reactions.
When methane burns, CH₄ + 2O₂ → CO₂ + 2H₂O, the bonds formed are 2 C=O double bonds in CO₂ and 4 O-H bonds in the two water molecules. This releases (2 × 805) + (4 × 464) = 3466 kJ/mol.
The energy to break a bond equals the energy released when that bond forms. A bond energy is the energy needed to break one mole of a particular covalent bond in the gaseous state, so the same values are used for bonds broken and bonds formed.
A water molecule has two O-H bonds, not one. Miscounting this is a common error in bond energy calculations; for example, 2H₂O has 4 O-H bonds in total, releasing 4 × 464 kJ/mol when they form.
Bond energy tables give positive values, because they are the energy needed to break bonds. The sign of the overall change comes from the formula energy in - energy out, so breaking bonds counts as energy in and forming bonds as energy out.
A positive energy change means the reaction is endothermic: more energy is needed to break bonds than is released when new bonds form, so energy is absorbed from the surroundings. A negative value means exothermic.