Energy is required to break chemical bonds, and energy is released when new bonds form. In bond energy calculations, you compare energy absorbed breaking reactant bonds against energy released forming product bonds: a negative result means the reaction is exothermic; a positive result means it is endothermic.
What is bond energy?
Bond energy (also called bond enthalpy) is the energy required to break one mole of a particular type of covalent bond in the gaseous state, measured in kJ/mol. It is always a positive value because breaking bonds always requires energy (an endothermic process).
For example:
- The C–H bond energy is 413 kJ/mol — this means 413 kJ of energy must be supplied to break one mole of C–H bonds
- The O=O bond energy is 498 kJ/mol — a double bond, so more energy is needed
Bond energies are average values because the energy needed to break a particular type of bond varies slightly depending on the surrounding atoms in the molecule. They are given as averages across many different molecules.
What are the rules about breaking and making bonds?
Two fundamental rules underpin every bond energy calculation:
| Process | Energy change | Name |
|---|---|---|
| Breaking bonds | Energy absorbed | Endothermic |
| Making bonds | Energy released | Exothermic |
A helpful way to remember: Break In, Make Out (energy goes in when bonds are broken; energy comes out when bonds are made).
The overall energy change for a reaction is:
ΔH = Σ(energy in to break reactant bonds) − Σ(energy out from forming product bonds)
- If ΔH is negative: more energy is released (making bonds) than absorbed (breaking bonds) → the reaction is exothermic
- If ΔH is positive: more energy is absorbed than released → the reaction is endothermic
Useful bond energies for GCSE calculations
| Bond | Bond energy (kJ/mol) |
|---|---|
| H–H | 436 |
| C–H | 413 |
| C–C | 347 |
| C=C | 614 |
| C–O | 360 |
| C=O | 805 |
| O=O | 498 |
| O–H | 463 |
| Cl–Cl | 243 |
| H–Cl | 432 |
| N≡N | 945 |
| N–H | 391 |
These are the values most commonly given in GCSE exam questions; you will not be expected to recall them from memory.
Worked example 1: is the reaction of hydrogen and chlorine exothermic or endothermic?
Reaction: H₂ + Cl₂ → 2HCl
Step 1 — Draw out the bonds in the reactants:
- 1 molecule of H₂ contains: 1 × H–H bond
- 1 molecule of Cl₂ contains: 1 × Cl–Cl bond
Step 2 — Calculate energy needed to break all reactant bonds:
| Bond broken | Number | Bond energy (kJ/mol) | Energy in (kJ) |
|---|---|---|---|
| H–H | 1 | 436 | 436 |
| Cl–Cl | 1 | 243 | 243 |
| Total energy in | 679 |
Step 3 — Draw out the bonds in the products:
- 2 molecules of HCl each contain: 1 × H–Cl bond → total 2 × H–Cl bonds
Step 4 — Calculate energy released forming all product bonds:
| Bond formed | Number | Bond energy (kJ/mol) | Energy out (kJ) |
|---|---|---|---|
| H–Cl | 2 | 432 | 864 |
| Total energy out | 864 |
Step 5 — Calculate ΔH:
ΔH = energy in − energy out = 679 − 864 = −185 kJ/mol
Conclusion: ΔH is negative → the reaction is exothermic. More energy is released making H–Cl bonds than was needed to break the H–H and Cl–Cl bonds.
Worked example 2: burning methane (combustion)
Reaction: CH₄ + 2O₂ → CO₂ + 2H₂O
Bonds in reactants:
- 1 × CH₄: contains 4 × C–H bonds
- 2 × O₂: contains 2 × O=O bonds
Energy in (breaking bonds):
- 4 × C–H: 4 × 413 = 1,652 kJ
- 2 × O=O: 2 × 498 = 996 kJ
- Total energy in = 1,652 + 996 = 2,648 kJ
Bonds in products:
- 1 × CO₂: contains 2 × C=O bonds
- 2 × H₂O: contains 2 × O–H bonds each → 4 × O–H bonds total
Energy out (making bonds):
- 2 × C=O: 2 × 805 = 1,610 kJ
- 4 × O–H: 4 × 463 = 1,852 kJ
- Total energy out = 1,610 + 1,852 = 3,462 kJ
ΔH = 2,648 − 3,462 = −814 kJ/mol
The combustion of methane is strongly exothermic, which is why natural gas is a useful fuel.
Why are bond energy calculations approximate?
Bond energies are average values — the actual energy of a C–H bond in methane is slightly different from the energy of a C–H bond in ethanol, because the surrounding atoms affect the electron density in the bond. When you use average values from a table, you are making an approximation.
The calculated ΔH from bond energies is therefore an estimate, not an exact value. The more accurate value can only be determined experimentally using calorimetry. Despite this limitation, bond energy calculations are useful for predicting whether a reaction will be exothermic or endothermic and for estimating the magnitude of the energy change.
Frequently asked questions
How do I know how many of each bond to break or make?
Read the balanced equation carefully and count every bond in every molecule on both sides. It helps to draw out the full structural formula of each molecule (showing all atoms and bonds), then count each type of bond. For example, in CH₄ there are 4 C–H bonds; in CO₂ there are 2 C=O double bonds; in H₂O there are 2 O–H bonds. Do not rely on the molecular formula alone — always draw or visualise the structure to avoid missing bonds.
Can a bond energy calculation give a positive ΔH, indicating an endothermic reaction?
Yes. If the energy needed to break bonds in the reactants is greater than the energy released when bonds form in the products, then ΔH = (energy in) − (energy out) is positive, which means the reaction absorbs energy from the surroundings — it is endothermic. The temperature of the reaction mixture falls. Photosynthesis is a biological example: the energy needed to break the bonds of water and carbon dioxide exceeds the energy released when making glucose bonds, and the deficit is supplied by absorbing light energy.
Why is bond breaking endothermic and bond making exothermic?
Covalent bonds form because the shared electrons are attracted to both nuclei — this arrangement is more stable (lower energy) than the atoms being separate. To break a bond, you must supply energy to separate the atoms back to a higher-energy, unbound state. Conversely, when two atoms come together and form a bond, the system moves to a lower energy state, and the excess energy is released to the surroundings. This is why breaking bonds always requires energy input and making bonds always releases energy.
What is the relationship between bond energy calculations and energy profile diagrams?
Both represent the same energy changes in a reaction. On an energy profile diagram, the activation energy (height of the peak above the reactants) corresponds to the energy needed to break reactant bonds at the transition state. The overall energy change (ΔH) is the difference between reactant and product energy levels — the same value you calculate from bond energies. A catalyst lowers the peak (reduces activation energy) but does not change the overall ΔH — the reactant and product bond energies are unchanged.
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