📘 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).
The Calculation
\( \Delta H = \sum(\text{bonds broken}) - \sum(\text{bonds formed}) \)
Breaking = positive (endo) | Forming = negative (exo)
Key Bond Enthalpies (from Data Booklet)
| Bond | kJ mol⁻¹ | Bond | kJ mol⁻¹ |
|---|---|---|---|
| H–H | 436 | C–C | 346 |
| C–H | 414 | C=C | 614 |
| O–H | 463 | O=O | 498 |
| C=O | 804 | N≡N | 945 |
⚠️ 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.
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