Bond energy
Section: 5 Chemical Energetics | Syllabus: Cambridge AS Level Chemistry (9701)
Bond Breaking and Bond Making
Enthalpy changes in reactions arise from the breaking and forming of chemical bonds:
- Bond breaking is always endothermic - energy must be supplied to overcome the attraction between atoms. ΔH is positive.
- Bond forming is always exothermic - energy is released as atoms come together and form new bonds. ΔH is negative.
- The overall ΔH of a reaction = energy put in to break bonds + energy released on forming bonds.
ΔH⦵ᵣ = Σ (bond energies of bonds broken) − Σ (bond energies of bonds formed)
If more energy is released forming bonds than is used breaking them → overall exothermic (ΔH negative).
Figure 5.6: Bond Breaking (Endothermic) vs Bond Forming (Exothermic)
(Two mini energy diagrams side by side. Left: bond breaking - two atoms separate, enthalpy level rises, ΔH positive, label "energy input required". Right: bond forming - two atoms come together, enthalpy level falls, ΔH negative, label "energy released". Show the atoms as spheres with a bond line between them.)
Bond Energy
Bond Energy (Mean Bond Enthalpy) The energy required to break one mole of a particular covalent bond in the gaseous state, to give gaseous atoms. Units: kJ mol⁻¹. Bond breaking is always endothermic (positive ΔH).Exact vs Average Bond Energies
- For diatomic molecules (H₂, Cl₂, HCl, N₂, O₂), the bond energy is an exact value - there is only one bond of that type.
- For bonds in more complex molecules (e.g. C–H bonds), the bond energy varies slightly depending on the molecular environment. The values quoted are mean (average) bond energies - averaged across many different compounds.
- Using mean bond energies gives approximate (not exact) values of ΔH. This is why bond energy calculations sometimes differ from measured values.
| Bond | Bond energy / kJ mol⁻¹ | Bond | Bond energy / kJ mol⁻¹ |
|---|---|---|---|
| H–H | 436 | C–H | 412 |
| Cl–Cl | 242 | C–C | 348 |
| H–Cl | 431 | C=C | 612 |
| O=O | 496 | C≡C | 837 |
| H–O | 463 | C=O | 743 |
| N≡N | 944 | C–O | 360 |
Higher bond energy = stronger bond = shorter bond length. Triple bonds are stronger and shorter than double bonds, which are stronger and shorter than single bonds (e.g. C≡C > C=C > C–C in energy and inversely in length).
Calculating ΔH from Bond Energies
Always draw the full structural formula first to identify every bond being broken and formed. Use fractions if necessary to balance atoms.
- Step 1: Draw the full structural formulae of reactants and products.
- Step 2: List all bonds broken in the reactants with their energies (positive values).
- Step 3: List all bonds formed in the products with their energies (positive values).
- Step 4: ΔH = Σ(broken) − Σ(formed)
Bonds broken (reactants):
4 × C–H = 4 × 412 = 1648 kJ
2 × O=O = 2 × 496 = 992 kJ
Total energy in = 1648 + 992 = 2640 kJ
Bonds formed (products):
2 × C=O (in CO₂) = 2 × 743 = 1486 kJ
4 × O–H (in 2H₂O) = 4 × 463 = 1852 kJ
Total energy out = 1486 + 1852 = 3338 kJ
ΔH = 2640 − 3338 = −698 kJ mol⁻¹
(Note: CO₂ has two C=O double bonds per molecule; H₂O has two O–H bonds per molecule.)
ΔH for formation of HCl: H₂ + Cl₂ → 2HClBonds broken: 1 × H–H (436) + 1 × Cl–Cl (242) = 678 kJ
Bonds formed: 2 × H–Cl (431) = 862 kJ
ΔH = 678 − 862 = −184 kJ mol⁻¹
Figure 5.7: Bond Energy Calculation - Energy Level Diagram for CH₄ Combustion
(Three-level enthalpy diagram: top level = gaseous atoms (all bonds broken, highest energy). From reactants CH₄ + 2O₂, arrow going up labelled "+2640 kJ (bonds broken)". From gaseous atoms, arrow going down to products CO₂ + 2H₂O labelled "−3338 kJ (bonds formed)". Net arrow from reactants to products labelled ΔH = −698 kJ mol⁻¹.)
Exam TipA common error is forgetting to multiply bond energies by the number of bonds. Count bonds carefully from the full structural formula - e.g. CH₄ has exactly 4 C–H bonds, and 2H₂O has 4 O–H bonds in total. Also note that CO₂ has two C=O bonds (O=C=O).
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