A bond enthalpy (or bond energy) is the energy required to break one mole of a particular covalent bond in the gaseous state, averaged across many different compounds.
Key Equation
\( \Delta H = \sum (\text{bonds broken}) - \sum (\text{bonds formed}) \)
Breaking bonds is endothermic (+). Making bonds is exothermic (−).
Worked Example
Example: Combustion of Methane
CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)
Bonds broken:
4 × C–H = 4 × 414 = 1656 kJ
2 × O=O = 2 × 498 = 996 kJ
Total = 2652 kJ
Bonds formed:
2 × C=O = 2 × 804 = 1608 kJ
4 × O–H = 4 × 463 = 1852 kJ
Total = 3460 kJ
ΔH = 2652 − 3460 = −808 kJ mol⁻¹
The literature value is −890 kJ mol⁻¹. The discrepancy arises because mean bond enthalpies are averages, not exact values for this specific molecule.
Why "Mean" Bond Enthalpies?
The energy of a C–H bond varies slightly depending on the molecule it is in (e.g. CH₄ vs C₂H₆). The mean bond enthalpy is an average across many compounds. This is why bond enthalpy calculations give approximate answers.
Think About It
Bond enthalpy calculations only apply to reactions in the gaseous state. Why can't we use them for reactions involving liquids or solids?
Because intermolecular forces (which require energy to overcome during state changes) are not accounted for in bond enthalpy values — only covalent bonds within molecules are considered.