Bond Dissociation Energy
Bond dissociation energy (BDE), also called bond energy, is the enthalpy change (ΔH) required to homolytically cleave a specific covalent bond in a gaseous molecule, breaking it into two neutral radical fragments. It is always a positive value (endothermic process) and serves as a direct measure of bond strength — higher BDE means a stronger, more stable bond. BDE values are used to estimate the enthalpy of reactions using Hess's law: bonds broken (endothermic) minus bonds formed (exothermic) gives the overall reaction enthalpy.
Key Formula
ΔH_rxn ≈ Σ BDE(bonds broken) − Σ BDE(bonds formed)
LaTeX: \Delta H_{\text{rxn}} \approx \sum \text{BDE}_{\text{broken}} - \sum \text{BDE}_{\text{formed}}
| Symbol | Meaning | Unit |
|---|---|---|
| ΔH_rxn | Enthalpy change of the reaction | kJ/mol |
| BDE_broken | Bond dissociation energy of bonds cleaved | kJ/mol |
| BDE_formed | Bond dissociation energy of bonds formed | kJ/mol |
Worked Example
Problem
Estimate the enthalpy change for the combustion of methane: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g). Use BDE values: C–H = 413 kJ/mol, O=O = 498 kJ/mol, C=O = 799 kJ/mol, O–H = 463 kJ/mol.
Solution
Step 1: Identify bonds broken in reactants. - CH₄: 4 × C–H bonds = 4 × 413 = 1652 kJ/mol - 2O₂: 2 × O=O bonds = 2 × 498 = 996 kJ/mol - Total bonds broken = 1652 + 996 = 2648 kJ/mol Step 2: Identify bonds formed in products. - CO₂: 2 × C=O bonds = 2 × 799 = 1598 kJ/mol - 2H₂O: 4 × O–H bonds = 4 × 463 = 1852 kJ/mol - Total bonds formed = 1598 + 1852 = 3450 kJ/mol Step 3: ΔH_rxn = 2648 – 3450 = –802 kJ/mol.
Answer
ΔH_rxn ≈ –802 kJ/mol. The reaction is strongly exothermic, consistent with the known standard enthalpy of combustion of methane (–890 kJ/mol; difference due to average BDE values used).
Average Bond Dissociation Energies of Common Bonds
| Bond | Bond Order | BDE (kJ/mol) | Example |
|---|---|---|---|
| H–H | Single | 436 | H₂ |
| C–H | Single | 413 | Methane, CH₄ |
| C–C | Single | 347 | Ethane, C₂H₆ |
| C=C | Double | 614 | Ethene, C₂H₄ |
| C≡C | Triple | 839 | Ethyne, C₂H₂ |
| N≡N | Triple | 945 | Dinitrogen, N₂ |
| O=O | Double | 498 | Dioxygen, O₂ |
| O–H | Single | 463 | Water, H₂O |
Interactive Tools
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Related Terms
Lattice Energy
Lattice energy is the energy required to completely separate one mole of a solid ionic compound into its constituent gaseous ions at infinite separation, or equivalently the energy released when gaseous ions combine to form the ionic lattice. It is always endothermic when defined as the energy of separation (positive value), and is a key measure of the stability and strength of an ionic compound. Lattice energy increases with increasing ionic charge and decreasing ionic radius, following the trend predicted by Coulomb's law.
Sigma Bond
A sigma bond (σ bond) is the strongest type of covalent bond, formed by the direct head-on overlap of atomic orbitals along the internuclear axis. It is the first bond formed between two atoms in any covalent bond and allows free rotation around the bond axis. Sigma bonds are present in all single, double, and triple bonds and are responsible for the overall framework and shape of molecules.
Pi Bond
A pi bond (π bond) is a covalent bond formed by the lateral (side-by-side) overlap of unhybridised p orbitals above and below the internuclear axis. Pi bonds are always formed in addition to an existing sigma bond, making up the second bond in a double bond and the second and third bonds in a triple bond. Unlike sigma bonds, pi bonds restrict rotation around the bond axis, which is critical for cis-trans isomerism in alkenes.
From Latin "dissociare" (to separate, disjoin) + "energia" (energy). The concept was formalised in thermochemistry in the early 20th century. The distinction between "bond dissociation energy" (specific bond in a specific molecule) and "mean bond energy" (average over many molecules) was clarified by Linus Pauling.