Bond Enthalpies

📘 Definition

Average bond enthalpy = the energy required to break one mole of a specific covalent bond in the gaseous state, averaged over a range of compounds. Always positive (endothermic).

\( \Delta H = \sum(\text{bonds broken}) - \sum(\text{bonds formed}) \)

Breaking = positive (endo) | Forming = negative (exo)

⚠️ Limitations. Why Answers are Approximate

• Bond enthalpies are averages. The actual C–H bond enthalpy varies between CH₄, C₂H₆, etc.
• Data applies to the gaseous state only. Ignores energy for changes of state
• Resonance structures (e.g. Benzene) make average values even less accurate

🔑 N≡N Insight

The triple bond in N₂ (945 kJ mol⁻¹) is one of the strongest bonds. This explains why nitrogen gas is so unreactive and why the Haber process needs extreme conditions.

📐 Worked Example: Using Bond Enthalpies

Calculate ΔH for: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)

Bonds broken (endothermic, positive):
4 × C-H = 4 × 414 = 1656 kJ
2 × O=O = 2 × 498 = 996 kJ
Total = 2652 kJ

Bonds formed (exothermic, negative):
2 × C=O = 2 × 804 = 1608 kJ
4 × O-H = 4 × 463 = 1852 kJ
Total = 3460 kJ

ΔH = bonds broken - bonds formed = 2652 - 3460 = -808 kJ mol^-1

Tip: A negative answer confirms the reaction is exothermic, which makes sense for combustion.

⚠️ Examiner Trap: Average Bond Enthalpies

Bond enthalpy values in the data booklet are averages taken across many different molecules. The actual bond enthalpy varies depending on the molecular environment. This means calculations using average bond enthalpies give approximate answers, not exact values.