Bond Enthalpies

Full Notes

Bond enthalpies have been outlined with more background theory here.

This page is just what you need to know for AQA A-level Chemistry :)

Understanding Bond Enthalpy

Bond enthalpy (bond dissociation energy) is defined as: "The energy required to break one mole of a particular bond in the gaseous state."

• Breaking bonds requires energy → endothermic (+ΔH).
• Making bonds releases energy → exothermic (-ΔH).

For Example:
Breaking H–H: H₂(g) → 2H(g), ΔH = +436 kJ mol⁻¹
Forming H–H: 2H(g) → H₂(g), ΔH = −436 kJ mol⁻¹

Mean Bond Enthalpy

Mean bond enthalpy is the average energy to break a bond, based on values across different compounds.

For Example:
C–H bond mean enthalpy ≈ +412 kJ mol⁻¹, but the actual enthalpy varies with environment in different molecules.

This means calculations using mean bond enthalpies to find enthalpy changes in reactions won’t be as accurate as those calculated using experimental data (such as from calorimetry).

Calculating Enthalpy Change Using Bond Enthalpies

The enthalpy change of reaction can be estimated as:

• Bonds broken (reactants): energy absorbed (positive ΔH).
• Bonds formed (products): energy released (negative ΔH).

Matt’s exam tip

Remember bond enthalpies are for gaseous states. If substances are liquids or solids, you may need to include enthalpy of vaporisation or fusion before applying bond enthalpies.

Ex

Worked Example: Combustion of Methane

Reaction: CH₄ + 2O₂ → CO₂ + 2H₂O

Bond enthalpies:

• C–H = +412 kJ·mol⁻¹
• O=O = +498 kJ·mol⁻¹
• C=O = +805 kJ·mol⁻¹
• O–H = +463 kJ·mol⁻¹

1. Bonds Broken (Reactants)
   • CH₄: 4 × C–H = 4 × 412 = 1648 kJ
   • O₂: 2 × O=O = 2 × 498 = 996 kJ
   • Total broken = 1648 + 996 = 2644 kJ
2. Bonds Formed (Products)
   • CO₂: 2 × C=O = 2 × 805 = 1610 kJ
   • H₂O: 4 × O–H = 4 × 463 = 1852 kJ
   • Total formed = 1610 + 1852 = 3462 kJ
3. Calculate ΔH
   ΔH = (bonds broken) − (bonds formed)
   ΔH = 2644 − 3462 = −818 kJ·mol⁻¹

ΔH = −818 kJ·mol⁻¹ (exothermic)

Accuracy of Bond Enthalpy Calculations

Experimental enthalpy values (e.g., from calorimetry) are usually more accurate than those calculated using average bond ethanlpies because:

• Average bond enthalpy values are approximations.
• We have to assume all substances are in gaseous state (not always true in real conditions).

Summary

• Bond enthalpy = energy to break 1 mol of a bond in the gaseous state.
• Breaking bonds is endothermic, forming bonds is exothermic.
• ΔH = Σ bonds broken − Σ bonds formed.
• Bond enthalpy calculations are approximate due to mean values and gaseous assumption.