The energy change of reactions (HT only)

What you'll learn

Why chemical reactions involve breaking bonds and forming bonds.
How to use bond energies to calculate an overall energy change.
How to decide whether a reaction is exothermic or endothermic from your calculation.
How to avoid the common Higher Tier mistakes with signs and bond counting.

The starting idea: reactions rearrange atoms

A chemical reaction happens when atoms are rearranged. The reactants are the starting substances, and the products are the substances made.

To rearrange atoms, old bonds in the reactants must be broken, and new bonds in the products must be made.

A chemical bond is a force of attraction holding particles together. In this topic you mostly deal with covalent bonds, where atoms share pairs of electrons in molecules.

Key Idea

Breaking absorbs, forming releases

Energy must be supplied to break bonds in the reactants. Energy is released when new bonds form in the products.

Diagram showing bond breaking absorbing energy, bond forming releasing energy, and how this links to exothermic and endothermic reactions

Exothermic and endothermic reactions

The surroundings are everything outside the reacting chemicals, such as the beaker, thermometer, water, and air around the reaction.

An exothermic reaction transfers energy to the surroundings. The surroundings usually get warmer.

An endothermic reaction takes in energy from the surroundings. The surroundings usually get colder.

Definition

Overall energy change

The overall energy change is the difference between the total energy absorbed when bonds are broken and the total energy released when bonds are formed. It is often represented as $ΔH\Delta H$ .

The overall energy change depends on which side “wins”:

If bond forming releases more energy than bond breaking needs, the reaction is exothermic.
If bond breaking needs more energy than bond forming releases, the reaction is endothermic.

Bond energy

Definition

Bond energy

A bond energy is the energy needed to break one mole of a particular type of covalent bond in gaseous molecules. It is measured in kilojoules per mole (kJ/mol). GCSE questions usually give you average bond energies.

For example, a C-H bond has a different bond energy from an O-H bond. A double bond such as O=O also has its own bond energy; do not treat it as two ordinary single O-O bonds unless the question tells you to.

Common Mistake

Bond-energy answers are estimates

Bond energies are average values, often based on gaseous molecules. Your calculated value may not exactly match an energy change measured in a practical using liquids, solutions, or imperfect insulation.

The Higher Tier calculation

For GCSE Chemistry Higher Tier, you need to calculate the energy transferred using bond energies supplied in the question.

The key equation is:

ΔH=Ebroken−Eformed\Delta H = E_{\text{broken}} - E_{\text{formed}}

where:

$EbrokenE_{\text{broken}}$ is the total energy needed to break bonds in the reactants.
$EformedE_{\text{formed}}$ is the total energy released when bonds form in the products.

So:

If $ΔH<0\Delta H < 0$ , the reaction is exothermic.
If $ΔH>0\Delta H > 0$ , the reaction is endothermic.

Common Mistake

Swapping the subtraction

Use bonds broken − bonds formed. If you calculate it the other way round, your sign will be wrong, so exothermic and endothermic reactions will be reversed.

Counting bonds accurately

Before you calculate anything, you must count how many of each bond type is broken and formed.

Look carefully at the balanced equation and any displayed formulae.

For example:

H₂ has one H-H bond.
Cl₂ has one Cl-Cl bond.
O₂ has one O=O bond.
CH₄ has four C-H bonds.
CO₂ has two C=O bonds.
H₂O has two O-H bonds, so 2H₂O has four O-H bonds.

Tip

Use the big number in front

A coefficient multiplies the whole molecule. So 2H₂O means there are two water molecules, each with two O-H bonds, giving four O-H bonds in total.

Example

Calculating the energy change for making hydrogen chloride

Hydrogen reacts with chlorine:

H₂(g) + Cl₂(g) → 2HCl(g)

Bond energies supplied:

H-H: 436 kJ/mol
Cl-Cl: 243 kJ/mol
H-Cl: 432 kJ/mol

Count the bonds broken in the reactants: one H-H bond and one Cl-Cl bond.
Calculate the energy needed to break these bonds:

436\ \text{kJ/mol} + 243\ \text{kJ/mol} = 679\ \text{kJ/mol}

Count the bonds formed in the products: 2HCl contains two H-Cl bonds.
Calculate the energy released when these bonds form:

\times 432\ \text{kJ/mol} = 864\ \text{kJ/mol}

Subtract using the GCSE bond-energy equation:

\Delta H = 679\ \text{kJ/mol} - 864\ \text{kJ/mol} = -185\ \text{kJ/mol}

Interpret the sign: the value is negative, so the reaction is exothermic. More energy is released forming H-Cl bonds than is needed to break the H-H and Cl-Cl bonds.

A full calculation with more bond types

Longer molecules are not harder in principle — you just need to be systematic.

Use this method:

Write or check the balanced equation.
Count all bonds broken in the reactants.
Multiply each number of bonds by its bond energy.
Add to get the total energy absorbed.
Count all bonds formed in the products.
Multiply each number of bonds by its bond energy.
Add to get the total energy released.
Calculate $ΔH=Ebroken−Eformed\Delta H = E_{\text{broken}} - E_{\text{formed}}$ .
Use the sign to decide exothermic or endothermic.

Example

Calculating the energy change for methane combustion

Methane burns in oxygen:

CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)

Bond energies supplied:

C-H: 413 kJ/mol
O=O: 498 kJ/mol
C=O: 805 kJ/mol
O-H: 464 kJ/mol

Count the bonds broken in the reactants: CH₄ has four C-H bonds, and 2O₂ has two O=O bonds.
Calculate the total energy needed to break bonds:

\left(4 \times 413\ \text{kJ/mol}\right) + \left(2 \times 498\ \text{kJ/mol}\right) = 2648\ \text{kJ/mol}

Count the bonds formed in the products: CO₂ has two C=O bonds, and 2H₂O has four O-H bonds.
Calculate the total energy released forming bonds:

\left(2 \times 805\ \text{kJ/mol}\right) + \left(4 \times 464\ \text{kJ/mol}\right) = 3466\ \text{kJ/mol}

Calculate the overall energy change:

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

Interpret the result: the reaction is exothermic because $ΔH\Delta H$ is negative. For the equation as written, 818 kJ/mol more energy is released in forming product bonds than is needed to break reactant bonds.

What your answer means

A negative answer does not mean “negative energy exists”. It means energy has been transferred from the reaction to the surroundings overall.

A positive answer means energy has been taken in from the surroundings overall.

Common Mistake

Saying bond breaking releases energy

Bond breaking always requires energy. The energy released in a reaction comes from forming new bonds, not from breaking old ones.

If a question asks for the energy change “per mole” of a particular substance, check the balanced equation. For example, if your equation makes 2 mol of product but the question wants the energy change per mole of that product, you would divide the reaction energy change by 2.

Exam technique

In the exam

Balance the equation first, then count bonds using the coefficients.
Use bonds broken − bonds formed, and keep the positive or negative sign.
If the answer is negative, write exothermic; if it is positive, write endothermic, and link this to whether bond forming released more or less energy than bond breaking required.

Self review

Check yourself

Why does breaking a chemical bond require energy rather than releasing it?
In 2H₂O(g), how many O-H bonds are formed?
What does a negative value of $ΔH\Delta H$ tell you about the reaction?

The energy change of reactions (HT only)

What you'll learn

Why chemical reactions involve breaking bonds and forming bonds.
How to use bond energies to calculate an overall energy change.
How to decide whether a reaction is exothermic or endothermic from your calculation.
How to avoid the common Higher Tier mistakes with signs and bond counting.

The starting idea: reactions rearrange atoms

A chemical reaction happens when atoms are rearranged. The reactants are the starting substances, and the products are the substances made.

To rearrange atoms, old bonds in the reactants must be broken, and new bonds in the products must be made.

A chemical bond is a force of attraction holding particles together. In this topic you mostly deal with covalent bonds, where atoms share pairs of electrons in molecules.

Key Idea

Breaking absorbs, forming releases

Energy must be supplied to break bonds in the reactants. Energy is released when new bonds form in the products.

Diagram showing bond breaking absorbing energy, bond forming releasing energy, and how this links to exothermic and endothermic reactions

Exothermic and endothermic reactions

The surroundings are everything outside the reacting chemicals, such as the beaker, thermometer, water, and air around the reaction.

An exothermic reaction transfers energy to the surroundings. The surroundings usually get warmer.

An endothermic reaction takes in energy from the surroundings. The surroundings usually get colder.

Definition

Overall energy change

The overall energy change is the difference between the total energy absorbed when bonds are broken and the total energy released when bonds are formed. It is often represented as $ΔH\Delta H$ .

The overall energy change depends on which side “wins”:

If bond forming releases more energy than bond breaking needs, the reaction is exothermic.
If bond breaking needs more energy than bond forming releases, the reaction is endothermic.

Bond energy

Definition

Bond energy

Common Mistake

Bond-energy answers are estimates

The Higher Tier calculation

For GCSE Chemistry Higher Tier, you need to calculate the energy transferred using bond energies supplied in the question.

The key equation is:

ΔH=Ebroken−Eformed\Delta H = E_{\text{broken}} - E_{\text{formed}}

where:

$EbrokenE_{\text{broken}}$ is the total energy needed to break bonds in the reactants.
$EformedE_{\text{formed}}$ is the total energy released when bonds form in the products.

So:

If $ΔH<0\Delta H < 0$ , the reaction is exothermic.
If $ΔH>0\Delta H > 0$ , the reaction is endothermic.

Common Mistake

Swapping the subtraction

Use bonds broken − bonds formed. If you calculate it the other way round, your sign will be wrong, so exothermic and endothermic reactions will be reversed.

Counting bonds accurately

Before you calculate anything, you must count how many of each bond type is broken and formed.

Look carefully at the balanced equation and any displayed formulae.

For example:

H₂ has one H-H bond.
Cl₂ has one Cl-Cl bond.
O₂ has one O=O bond.
CH₄ has four C-H bonds.
CO₂ has two C=O bonds.
H₂O has two O-H bonds, so 2H₂O has four O-H bonds.

Tip

Use the big number in front

A coefficient multiplies the whole molecule. So 2H₂O means there are two water molecules, each with two O-H bonds, giving four O-H bonds in total.

Example

Calculating the energy change for making hydrogen chloride

Hydrogen reacts with chlorine:

H₂(g) + Cl₂(g) → 2HCl(g)

Bond energies supplied:

H-H: 436 kJ/mol
Cl-Cl: 243 kJ/mol
H-Cl: 432 kJ/mol

Count the bonds broken in the reactants: one H-H bond and one Cl-Cl bond.
Calculate the energy needed to break these bonds:

436\ \text{kJ/mol} + 243\ \text{kJ/mol} = 679\ \text{kJ/mol}

Count the bonds formed in the products: 2HCl contains two H-Cl bonds.
Calculate the energy released when these bonds form:

\times 432\ \text{kJ/mol} = 864\ \text{kJ/mol}

Subtract using the GCSE bond-energy equation:

\Delta H = 679\ \text{kJ/mol} - 864\ \text{kJ/mol} = -185\ \text{kJ/mol}

Interpret the sign: the value is negative, so the reaction is exothermic. More energy is released forming H-Cl bonds than is needed to break the H-H and Cl-Cl bonds.

A full calculation with more bond types

Longer molecules are not harder in principle — you just need to be systematic.

Use this method:

Write or check the balanced equation.
Count all bonds broken in the reactants.
Multiply each number of bonds by its bond energy.
Add to get the total energy absorbed.
Count all bonds formed in the products.
Multiply each number of bonds by its bond energy.
Add to get the total energy released.
Calculate $ΔH=Ebroken−Eformed\Delta H = E_{\text{broken}} - E_{\text{formed}}$ .
Use the sign to decide exothermic or endothermic.

Example

Calculating the energy change for methane combustion

Methane burns in oxygen:

CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)

Bond energies supplied:

C-H: 413 kJ/mol
O=O: 498 kJ/mol
C=O: 805 kJ/mol
O-H: 464 kJ/mol

Count the bonds broken in the reactants: CH₄ has four C-H bonds, and 2O₂ has two O=O bonds.
Calculate the total energy needed to break bonds:

\left(4 \times 413\ \text{kJ/mol}\right) + \left(2 \times 498\ \text{kJ/mol}\right) = 2648\ \text{kJ/mol}

Count the bonds formed in the products: CO₂ has two C=O bonds, and 2H₂O has four O-H bonds.
Calculate the total energy released forming bonds:

\left(2 \times 805\ \text{kJ/mol}\right) + \left(4 \times 464\ \text{kJ/mol}\right) = 3466\ \text{kJ/mol}

Calculate the overall energy change:

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

Interpret the result: the reaction is exothermic because $ΔH\Delta H$ is negative. For the equation as written, 818 kJ/mol more energy is released in forming product bonds than is needed to break reactant bonds.

What your answer means

A negative answer does not mean “negative energy exists”. It means energy has been transferred from the reaction to the surroundings overall.

A positive answer means energy has been taken in from the surroundings overall.

Common Mistake

Saying bond breaking releases energy

Bond breaking always requires energy. The energy released in a reaction comes from forming new bonds, not from breaking old ones.

Exam technique

In the exam

Balance the equation first, then count bonds using the coefficients.
Use bonds broken − bonds formed, and keep the positive or negative sign.
If the answer is negative, write exothermic; if it is positive, write endothermic, and link this to whether bond forming released more or less energy than bond breaking required.

Self review

Check yourself

Why does breaking a chemical bond require energy rather than releasing it?
In 2H₂O(g), how many O-H bonds are formed?
What does a negative value of $ΔH\Delta H$ tell you about the reaction?

The energy change of reactions (HT only)

What you'll learn

The starting idea: reactions rearrange atoms

Exothermic and endothermic reactions

Bond energy

The Higher Tier calculation

Counting bonds accurately

A full calculation with more bond types

What your answer means

The energy change of reactions (HT only) Revision Guide

The energy change of reactions (HT only)

What you'll learn

The starting idea: reactions rearrange atoms

Exothermic and endothermic reactions

Bond energy

The Higher Tier calculation

Counting bonds accurately

A full calculation with more bond types

What your answer means

The energy change of reactions (HT only) Revision Guide