1.6.7 Bond Enthalpies

What are Bond Enthalpies?

Bond enthalpy (also known as bond dissociation enthalpy) is the amount of energy required to break one mole of a specific covalent bond in a molecule, in the gaseous state. This process is always endothermic because energy is required to break bonds, and therefore the enthalpy change ( $ΔH$ ) is positive.

Mean Bond Enthalpy

The mean bond enthalpy is the average energy needed to break one mole of a specific type of bond across a range of compounds where that bond exists. Since the strength of a bond can vary slightly depending on its chemical environment, the "mean" value accounts for these variations.

Breaking bonds: Always requires energy, so bond breaking is endothermic and $ΔH$ is positive.
Forming bonds: Always releases energy, so bond formation is exothermic and $ΔH$ is negative.

General Trends

Stronger bonds require more energy to break and release more energy when formed.
Shorter bonds are typically stronger:
- Triple bonds are generally shorter and stronger than double bonds.
- Double bonds are shorter and stronger than single bonds.

Calculating Enthalpy Changes Using Mean Bond Enthalpies

The enthalpy change ( $ΔH$ ) of a reaction can be calculated using mean bond enthalpies by comparing the energy needed to break the bonds in the reactants to the energy released when new bonds form in the products.

The formula is:

ΔH = \sum (\text{mean bond enthalpies of bonds broken}) - \sum (\text{mean bond enthalpies of bonds formed})

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Example: For the reaction:

H_2(g) + Cl_2(g) \rightarrow 2HCl(g)

If the bond enthalpies are:

$H-H$ bond = 436 kJ/mol
$Cl-Cl$ bond = 243 kJ/mol
$H-Cl$ bond = 431 kJ/mol The calculation would be:

Step 1: Energy to break the bonds in the reactants:

436 + 243 = :highlight[679 \, \text{kJ/mol}]

Step 2: Energy released when forming new bonds in the products:

2 \times 431 = :highlight[862 \, \text{kJ/mol}]

Step 3: Overall enthalpy change:

ΔH = 679 - 862 = :highlight[-183 \, \text{kJ/mol}]

This is an exothermic reaction because $ΔH$ is negative.

Differences Between Mean Bond Enthalpy Calculations and Hess' Law

Calculations using mean bond enthalpies often differ from those done using Hess' Law for several reasons:

Mean bond enthalpy values are averages, and actual bond energies can vary depending on the chemical environment of the bond.
State of the substances: Mean bond enthalpy calculations assume all substances are in the gaseous state. However, real reactions may involve liquids or solids, where intermolecular forces also contribute to the enthalpy change.
Accuracy: Calculations based on enthalpies of formation or combustion are generally more accurate because they are experimentally determined for specific reactions under standard conditions, rather than using average bond energies.

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Summary

Bond enthalpies represent the energy required to break a bond, with mean bond enthalpy being an average value across different compounds.
Calculating $ΔH$ using bond enthalpies involves summing the energies of bonds broken and formed.
Mean bond enthalpy calculations are approximations and can differ from more accurate methods like Hess' Law due to averaging and differences in physical states.

Bond Enthalpies Simplified Revision Notes for A-Level AQA Chemistry