The energy change of reactions (HT only) — AQA GCSE Chemistry
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The energy change of reactions (HT only) explained
The energy change of reactions (HT only) covers 6 specification points, including Energy must be supplied to break bonds in the… and The energy needed to break bonds and the….
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Open a point below for its explanation and revision guidance.
During a chemical reaction: energy must be supplied to break bonds in the reactants energy is released when bonds in the products are formed.
Chemical reactions involve breaking the bonds in the reactants and making new bonds in the products. Breaking a bond requires an input of energy because the atoms are pulled apart against the attraction holding them together, so bond breaking is endothermic. Forming a bond releases energy as the atoms come together and the attraction pulls them into a lower-energy arrangement, so bond making is exothermic. The overall energy change is the balance of these two processes: total energy needed to break reactant bonds minus total energy released when product bonds form. If more energy is released than supplied, the reaction is exothermic overall; if more is supplied than released, it is endothermic. Bond energies, usually quoted in kJ/mol, allow this balance to be calculated.
The energy needed to break bonds and the energy released when bonds are formed can be calculated from bond energies.
Bond energies are average values, usually quoted in kJ/mol, for breaking one mole of a given bond in the gaseous state. To calculate the energy needed to break bonds in the reactants, identify every bond in the reactant molecules, multiply each bond energy by the number of that bond present, and add the values. To calculate the energy released when bonds form in the products, do the same for the product molecules. For example, in H₂ + Cl₂ → 2HCl, breaking H–H and Cl–Cl requires 436 + 243 = 679 kJ/mol, while forming two H–Cl bonds releases 2 × 431 = 862 kJ/mol. Bond breaking is endothermic and bond forming is exothermic, so the two totals are calculated separately before being combined.
The difference between the sum of the energy needed to break bonds in the reactants and the sum of the energy released when bonds in the products are formed is the overall energy change of the reaction.
The overall energy change of a reaction is found by subtracting the energy released when new bonds form from the energy needed to break existing bonds. In symbols, ΔH = energy needed to break bonds − energy released when bonds form. For H₂ + Cl₂ → 2HCl, breaking bonds needs 436 + 243 = 679 kJ/mol and forming bonds releases 2 × 431 = 862 kJ/mol, so ΔH = 679 − 862 = −183 kJ/mol. A negative result means the reaction is exothermic overall; a positive result means it is endothermic overall. The calculation must use the balanced equation so that the number of each bond is correct, and the sign of the answer must be interpreted, not just calculated.
In an exothermic reaction, the energy released from forming new bonds is greater than the energy needed to break existing bonds.
An exothermic reaction transfers energy to the surroundings, usually shown by a temperature rise. At the bond level, the energy released when new bonds form in the products is greater than the energy needed to break bonds in the reactants. For example, in the combustion of methane, breaking all the bonds in CH₄ and O₂ requires a certain amount of energy, but forming the bonds in CO₂ and H₂O releases more energy, so the overall energy change is negative and the reaction is exothermic. The excess energy is transferred to the surroundings. A reaction can still need an initial energy input to start, but that does not make it endothermic overall.
In an endothermic reaction, the energy needed to break existing bonds is greater than the energy released from forming new bonds.
An endothermic reaction takes in energy from the surroundings, often shown by a temperature fall. At the bond level, the energy needed to break existing bonds in the reactants is greater than the energy released when new bonds form in the products. For example, thermal decomposition of calcium carbonate requires continuous heating because breaking the bonds in CaCO₃ needs more energy than is released when the products form, so the overall energy change is positive. The surroundings supply the extra energy, which is why the temperature of the surroundings can fall. The sign of the calculated overall energy change confirms whether the reaction is endothermic.
Students should be able to calculate the energy transferred in chemical reactions using bond energies supplied.
Bond breaking is endothermic because energy must be supplied to overcome attractions, while bond making is exothermic because energy is released when new bonds form. To find the overall energy transferred, identify every bond broken in the reactants and every bond made in the products, using the supplied bond energies in kJ mol⁻¹. Total the energy absorbed for bonds broken and the energy released for bonds made. The overall energy change is energy absorbed minus energy released. A negative result means the reaction is exothermic overall; a positive result means it is endothermic. For example, in the reaction H₂ + Cl₂ → 2HCl, breaking one H–H bond and one Cl–Cl bond absorbs energy, while forming two H–Cl bonds releases energy. The difference gives the overall energy change.
Your focus
- Explain why bond breaking requires energy and bond formation releases energy.
- Calculate the overall energy change of a reaction from bond energies.
- Use the sign of the overall energy change to classify a reaction as exothermic or endothermic.
Show all 18 objectives
- Identify all bonds broken in the reactants and all bonds formed in the products from a balanced equation.
- Calculate the total energy needed to break bonds and the total energy released when bonds form using supplied bond energies.
- Explain why bond breaking requires energy input while bond formation releases energy.
- Calculate the overall energy change of a reaction from bond energies using the correct subtraction order.
- Interpret the sign of an overall energy change as showing an exothermic or endothermic reaction.
- Present a bond energy calculation with correct units, sign and clear working.
- Explain why an exothermic reaction releases energy in terms of bond breaking and bond formation.
- Use bond energy totals to show that the energy released when bonds form is greater than the energy needed to break bonds.
- Relate a negative overall energy change to an observable temperature rise in the surroundings.
- Explain why an endothermic reaction takes in energy in terms of bond breaking and bond formation.
- Use bond energy totals to show that the energy needed to break bonds is greater than the energy released when bonds form.
- Relate a positive overall energy change to an observable temperature fall or the need for continuous heating.
- Identify all bonds broken and made from a supplied equation or displayed structure.
- Calculate total energy absorbed and total energy released using supplied bond energies.
- Determine the overall energy change and classify the reaction as exothermic or endothermic.
The energy change of reactions (HT only) exam tips
Marking Points
- Bond breaking in reactants requires an input of energy and is endothermic.
- Bond formation in products releases energy to the surroundings and is exothermic.
- The overall energy change equals the energy needed to break bonds in the reactants minus the energy released when bonds form in the products.
- A negative overall energy change indicates an exothermic reaction; a positive value indicates an endothermic reaction.
- Bond energies are average values in kJ/mol and can be used to calculate the overall energy change.
- The same bonds may appear in reactants and products, so only the net difference contributes to the overall change.
- State that bond breaking is endothermic because energy must be supplied to overcome the attraction holding the atoms together.
- State that bond formation is exothermic because energy is released as new bonds form and the system becomes more stable.
- Use bond energies in kJ/mol as average values for breaking one mole of a specified bond in the gaseous state.
- Calculate the total energy needed to break bonds by multiplying each bond energy by the number of that bond in the reactants and adding the values.
- Calculate the total energy released when bonds form by multiplying each bond energy by the number of that bond in the products and adding the values.
- Show clear working that distinguishes the breaking total from the forming total before any subtraction is carried out.
- Calculate the total energy needed to break bonds in the reactants by adding the relevant bond energies.
- Calculate the total energy released when bonds form in the products by adding the relevant bond energies.
- Subtract the energy released when bonds form from the energy needed to break bonds to obtain the overall energy change.
- Interpret a negative overall energy change as exothermic and a positive overall energy change as endothermic.
- Include the correct sign and units, usually kJ/mol, with the final answer.
- Use the balanced equation to ensure the number of each bond broken and formed is correct.
- State that in an exothermic reaction the energy released when new bonds form is greater than the energy needed to break existing bonds.
- Link the greater energy release to a negative overall energy change, ΔH.
- Describe the transfer of excess energy to the surroundings, often observed as a temperature rise.
- Use a balanced equation and bond energies to show numerically that the forming total exceeds the breaking total.
- Distinguish the initial activation energy input from the overall energy change of the reaction.
- State that in an endothermic reaction the energy needed to break existing bonds is greater than the energy released when new bonds form.
- Link the greater energy input to a positive overall energy change, ΔH.
- Describe energy being taken in from the surroundings, often observed as a temperature fall.
- Use a balanced equation and bond energies to show numerically that the breaking total exceeds the forming total.
- Explain that continuous heating may be needed to supply energy for the reaction to continue.
- State that bond breaking is endothermic and bond making is exothermic.
- Identify all bonds broken in the reactants and all bonds made in the products from a displayed or supplied equation.
- Multiply each bond energy by the number of that bond present, using the supplied values in kJ mol⁻¹.
- Calculate total energy absorbed for bonds broken and total energy released for bonds made.
- Apply overall energy change = energy absorbed − energy released, and interpret the sign as exothermic or endothermic.
Examiner Tips
- 💡Write the calculation as a clear subtraction of total bond energies, showing the sum for bonds broken and the sum for bonds made.
- 💡Check the sign of your final answer against the reaction type: negative for exothermic, positive for endothermic.
- 💡Quote bond energy values with units of kJ/mol and keep significant figures consistent through the calculation.
- 💡Draw displayed formulae for all reactants and products before selecting bond energies, so no bond is missed or double-counted.
- 💡Write the breaking total and the forming total on separate lines with units of kJ/mol, then combine them in a clearly labelled final step.
- 💡Check that the equation is balanced before counting bonds, because coefficients change the number of bonds present.
- 💡Set out the calculation in two labelled stages: energy to break bonds, then energy released when bonds form, before subtracting.
- 💡State the final answer with its sign and units, and add a short sentence interpreting the sign as exothermic or endothermic.
- 💡If the question supplies bond energies in a table, copy each value accurately and show which bond it applies to.
- 💡Use the phrase energy released when bonds form is greater than energy needed to break bonds to make the comparison explicit.
- 💡Support the explanation with a calculated overall energy change that is negative, including units.
- 💡Link the bond-level explanation to an observable change such as a temperature rise of the surroundings.
- 💡State the comparison explicitly: energy needed to break bonds is greater than energy released when bonds form.
- 💡Support the explanation with a calculated overall energy change that is positive, including units.
- 💡Link the bond-level explanation to an observable change such as a temperature fall or the need for continuous heating.
- 💡Draw the displayed structures of reactants and products before counting bonds, so no bond is missed or double-counted.
- 💡Write a clear two-column layout: bonds broken with total absorbed, then bonds made with total released.
- 💡Check the final sign against the reaction context; combustion and neutralisation are normally exothermic overall.
Common Mistakes
- Stating that bond breaking releases energy: bond breaking always requires energy input, while bond formation releases energy.
- Subtracting in the wrong order, for example products minus reactants when the convention used is reactants minus products; keep one consistent order and interpret the sign.
- Assuming a reaction is exothermic simply because bonds form in the products, ignoring the energy needed to break the reactant bonds first.
- Adding bond energies for bonds that are not present in the equation; correct this by drawing displayed formulae for every reactant and product and counting each bond carefully.
- Using the number of molecules rather than the number of bonds; correct this by multiplying the bond energy by the number of that specific bond in the balanced equation.
- Treating bond formation as endothermic; correct this by remembering that forming bonds releases energy and only breaking bonds requires energy input.
- Subtracting in the wrong order, for example energy to break bonds minus energy released when bonds form reversed; correct this by always using energy needed to break bonds minus energy released when bonds form.
- Omitting the sign of the final answer; correct this by stating whether the value is positive or negative and linking that sign to endothermic or exothermic behaviour.
- Forgetting to multiply bond energies by coefficients in the balanced equation; correct this by counting bonds from displayed formulae that match the balanced equation.
- Claiming that exothermic reactions need no energy to break bonds; correct this by explaining that bonds must still be broken, but forming bonds releases more energy overall.
- Confusing a temperature rise in the surroundings with energy being created; correct this by describing energy transfer from the reaction to the surroundings.
- Assuming any reaction that starts with heating must be endothermic; correct this by comparing the total energy released when bonds form with the total energy needed to break bonds.
- Stating that endothermic reactions release energy because new bonds form; correct this by explaining that bond formation does release energy, but less than the energy needed to break the existing bonds.
- Ignoring the sign of the overall energy change; correct this by stating that a positive value indicates an endothermic reaction.
- Confusing a temperature fall in the surroundings with the reaction mixture becoming colder by creating coldness; correct this by describing energy transfer from the surroundings to the reaction.
- Adding bond energies for bonds made instead of subtracting them; correct by calculating energy absorbed minus energy released.
- Forgetting to multiply a bond energy by the number of identical bonds, such as two H–Cl bonds in 2HCl; correct by counting every bond in the equation.
- Reversing the sign convention, so an exothermic reaction is reported as positive; correct by checking that energy released greater than energy absorbed gives a negative overall value.