Energy changes: bond energies and calculating energy changes

What you'll learn

This revision guide covers how to calculate energy changes in chemical reactions using bond energies. You'll learn why breaking bonds requires energy input whilst forming bonds releases energy, and how to use these values to determine whether a reaction is exothermic or endothermic. These concepts are essential for AQA GCSE Chemistry Paper 1 and form the foundation for understanding energy transfers in all chemical reactions.

Key terms and definitions

Bond energy — the energy required to break one mole of a particular covalent bond in the gaseous state (measured in kJ/mol)

Endothermic reaction — a reaction that takes in energy from the surroundings, resulting in a temperature decrease and a positive energy change

Exothermic reaction — a reaction that transfers energy to the surroundings, resulting in a temperature increase and a negative energy change

Energy change of a reaction — the difference between the energy required to break bonds in the reactants and the energy released when forming bonds in the products

Activation energy — the minimum energy required for a reaction to occur, needed to start breaking bonds in the reactants

Bond breaking — an endothermic process that requires energy input to overcome the attractive forces between bonded atoms

Bond forming — an exothermic process that releases energy as new attractive forces are established between atoms

Core concepts

Understanding bond energies

Bond energies represent the strength of covalent bonds between atoms. When you see a bond energy value of 436 kJ/mol for H–H, this means breaking one mole of hydrogen-hydrogen bonds requires 436 kJ of energy.

Key points about bond energies:

• Bond energies are always quoted as positive values because breaking bonds always requires energy input
• The same amount of energy released when forming a bond equals the energy needed to break it
• Bond energies are measured for gaseous molecules only
• Different bonds have different strengths (e.g., C=C is stronger than C–C)
• Bond energies are average values because the same bond can have slightly different strengths in different molecules

Common bond energies you should recognise:

• C–H: approximately 413 kJ/mol
• O=O: approximately 498 kJ/mol
• C=O: approximately 805 kJ/mol
• H–H: approximately 436 kJ/mol
• O–H: approximately 464 kJ/mol
• C–C: approximately 347 kJ/mol

Bond breaking and bond forming

Every chemical reaction involves both breaking existing bonds and forming new ones. Understanding this two-step process is crucial for energy calculations.

Breaking bonds (endothermic):

When reactant molecules break apart, energy must be supplied to overcome the attractive forces holding atoms together. This energy input is always endothermic—it takes energy from the surroundings.

Example: Breaking O=O in oxygen gas requires 498 kJ/mol of energy input.

Forming bonds (exothermic):

When new bonds form in product molecules, energy is released as atoms become attracted to each other and form stable bonds. This energy release is always exothermic—it transfers energy to the surroundings.

Example: Forming O–H bonds in water releases 464 kJ/mol per bond.

Calculating energy changes using bond energies

The overall energy change of a reaction depends on the balance between energy required to break bonds and energy released when forming bonds.

The calculation method:

1. Calculate the total energy required to break all bonds in the reactants
2. Calculate the total energy released when forming all bonds in the products
3. Find the difference: Energy change = Energy required to break bonds − Energy released forming bonds

Interpreting the result:

• If energy change is positive: the reaction is endothermic (more energy needed to break bonds than released forming them)
• If energy change is negative: the reaction is exothermic (more energy released forming bonds than needed to break them)

Energy profiles and bond energies

Energy profile diagrams show how energy changes during a reaction. These diagrams illustrate:

For exothermic reactions:

• Products have lower energy than reactants
• The difference represents energy released to surroundings
• Overall energy change is negative
• Temperature of surroundings increases

For endothermic reactions:

• Products have higher energy than reactants
• The difference represents energy absorbed from surroundings
• Overall energy change is positive
• Temperature of surroundings decreases

Both types require activation energy to initiate bond breaking before the reaction proceeds.

Drawing and interpreting structural formulae for bond calculations

To calculate bond energies accurately, you must identify every bond in the molecules involved. Use displayed structural formulae showing all atoms and bonds.

Counting bonds correctly:

• Each line in a structural formula represents one bond
• Double bonds (=) count as one C=C or O=O bond, not two separate bonds
• Don't forget bonds to hydrogen atoms—these are often where mistakes occur
• Methane (CH₄) has four C–H bonds, not one
• Water (H₂O) has two O–H bonds

State symbols matter:

Bond energy calculations only apply to reactions where all substances are gaseous (g). If you're given state symbols showing liquids or solids, the calculation will be approximate because bond energies are defined for gaseous states only.

Using bond energy tables

In AQA GCSE Chemistry exams, bond energy values are always provided in questions—you don't need to memorise them. However, you must know how to use the tables correctly.

Reading the table:

• Bond energies are typically given in kJ/mol
• The table shows the bond type (e.g., C–H, O=O)
• Remember these are average values
• Use the exact values given in the question, not rounded figures until the final answer

Applying values to calculations:

• Identify which bonds are present in your molecules
• Count how many of each bond type
• Multiply the number of each bond by its bond energy
• Add up all the bond energies for all bonds being broken or formed

Worked examples

Example 1: Combustion of methane

Question: Calculate the energy change for the complete combustion of methane:

CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)

Bond energies: C–H = 413 kJ/mol, O=O = 498 kJ/mol, C=O = 805 kJ/mol, O–H = 464 kJ/mol

Solution:

Step 1: Identify bonds broken in reactants

• Methane (CH₄): 4 × C–H bonds
• Oxygen (2O₂): 2 × O=O bonds

Step 2: Calculate energy required to break bonds

• Breaking C–H bonds: 4 × 413 = 1652 kJ/mol
• Breaking O=O bonds: 2 × 498 = 996 kJ/mol
• Total energy in = 1652 + 996 = 2648 kJ/mol

Step 3: Identify bonds formed in products

• Carbon dioxide (CO₂): 2 × C=O bonds
• Water (2H₂O): 4 × O–H bonds (2 molecules × 2 bonds each)

Step 4: Calculate energy released forming bonds

• Forming C=O bonds: 2 × 805 = 1610 kJ/mol
• Forming O–H bonds: 4 × 464 = 1856 kJ/mol
• Total energy out = 1610 + 1856 = 3466 kJ/mol

Step 5: Calculate overall energy change Energy change = Energy in − Energy out Energy change = 2648 − 3466 = −818 kJ/mol

The negative value indicates this is an exothermic reaction releasing 818 kJ/mol.

Mark scheme points: Correct identification of all bonds (1 mark), correct calculation of energy in (1 mark), correct calculation of energy out (1 mark), correct overall energy change with sign (1 mark)

Example 2: Hydrogen and chlorine reaction

Question: Hydrogen reacts with chlorine to form hydrogen chloride:

H₂(g) + Cl₂(g) → 2HCl(g)

Use bond energies to calculate the energy change: H–H = 436 kJ/mol, Cl–Cl = 243 kJ/mol, H–Cl = 432 kJ/mol

Solution:

Bonds broken (reactants):

• 1 × H–H = 436 kJ/mol
• 1 × Cl–Cl = 243 kJ/mol
• Total energy required = 436 + 243 = 679 kJ/mol

Bonds formed (products):

• 2 × H–Cl = 2 × 432 = 864 kJ/mol

Energy change = 679 − 864 = −185 kJ/mol

This exothermic reaction releases 185 kJ/mol of energy.

Mark scheme points: All bonds correctly identified (1 mark), energy to break bonds = 679 kJ/mol (1 mark), energy released = 864 kJ/mol (1 mark), final answer −185 kJ/mol (1 mark)

Example 3: Decomposition of ammonia

Question: Calculate whether the decomposition of ammonia is endothermic or exothermic:

2NH₃(g) → N₂(g) + 3H₂(g)

Bond energies: N–H = 391 kJ/mol, N≡N = 945 kJ/mol, H–H = 436 kJ/mol

Solution:

Bonds broken in reactants:

• 2 molecules of NH₃, each with 3 × N–H bonds
• Total N–H bonds = 6
• Energy required = 6 × 391 = 2346 kJ/mol

Bonds formed in products:

• 1 × N≡N = 945 kJ/mol
• 3 × H–H = 3 × 436 = 1308 kJ/mol
• Total energy released = 945 + 1308 = 2253 kJ/mol

Energy change = 2346 − 2253 = +93 kJ/mol

The positive value indicates this is an endothermic reaction requiring 93 kJ/mol of energy input.

Mark scheme points: Correct count of 6 N–H bonds broken (1 mark), energy in = 2346 kJ/mol (1 mark), energy out = 2253 kJ/mol (1 mark), positive energy change showing endothermic (1 mark)

Common mistakes and how to avoid them

• Forgetting to multiply by the number of bonds present. If there are 4 C–H bonds in methane, you must multiply the C–H bond energy by 4, not just use the value once. Always count carefully using displayed formulae.

• Confusing which energy is 'in' and which is 'out'. Remember: breaking bonds requires energy IN (reactants), forming bonds releases energy OUT (products). A common error is reversing this in the final calculation.

• Getting the sign wrong on the final answer. Exothermic reactions should have negative energy changes (energy released), endothermic reactions have positive energy changes (energy absorbed). If your answer seems opposite to what you expect, check your subtraction.

• Missing hydrogen atoms when counting bonds. Hydrogen atoms are small and easy to miss in structural formulae. In water (H₂O), there are two O–H bonds; in methane (CH₄), there are four C–H bonds. Draw out the full displayed formula if needed.

• Not reading bond energy tables carefully. C–C, C=C, and C≡C are three different bonds with different energies. Similarly, C–O and C=O are different. Make sure you select the correct bond type from the table.

• Rounding too early in calculations. Keep full values throughout your working and only round the final answer. Rounding intermediate values can lead to significant errors in multi-step calculations.

Exam technique for "Energy changes: bond energies and calculating energy changes"

• Show all working clearly in separate steps. Marks are awarded for method even if your final answer is incorrect. Write out: bonds broken (with calculation), bonds formed (with calculation), then overall energy change. This structured approach typically earns 3-4 marks even with arithmetic errors.

• Include units (kJ/mol) and signs (+ or −) in your final answer. The sign indicates whether the reaction is exothermic or endothermic—this is often worth a mark. Leaving out the negative sign on an exothermic reaction loses marks.

• Command word 'Calculate' requires numerical working. You must show the calculation steps, not just state an answer. Command word 'Explain' requires you to link bond breaking/forming to energy changes and state whether the reaction is exothermic or endothermic with reasoning.

• Use the bond energies provided in the question. Don't use memorised values even if you know them—examiners sometimes use slightly different values to test your method. Always refer to the data given.

Quick revision summary

Bond energies measure the energy needed to break one mole of a specific covalent bond. Breaking bonds is endothermic (requires energy); forming bonds is exothermic (releases energy). To calculate energy changes: find the total energy needed to break all reactant bonds, subtract the total energy released forming product bonds. Negative results indicate exothermic reactions (more energy out than in); positive results indicate endothermic reactions (more energy in than out). Always show clear working, use displayed formulae to count bonds accurately, and include signs and units in answers.