Bond Enthalpy Formula
Estimate reaction enthalpy from bond energies
Understanding Bond Enthalpy
Bond enthalpy (also called bond energy or bond dissociation energy) is the energy required to break one mole of a particular type of bond in gaseous molecules under standard conditions, producing gaseous atoms or radicals. This fundamental thermochemical concept allows chemists to estimate reaction enthalpies (ΔHrxn) by considering the energy changes involved in breaking existing bonds in reactants and forming new bonds in products. The basic principle is elegant: breaking bonds requires energy input (endothermic, positive values) while forming bonds releases energy (exothermic, negative values).
Bond enthalpy values are typically reported as average values across many molecules because the exact energy required to break a specific bond depends on its molecular environment. For example, the C-H bond energy in methane (CH₄) differs slightly from that in ethane (C₂H₆) or benzene (C₆H₆) due to differences in neighboring atoms and molecular structure. Despite this limitation, average bond enthalpies provide remarkably useful estimates for predicting reaction energetics, especially when more precise thermodynamic data (like standard enthalpies of formation) are unavailable.
The bond enthalpy approach is particularly valuable in organic chemistry where countless reaction pathways exist, making it impractical to measure ΔHrxn experimentally for every possible transformation. By calculating ΔHrxn = Σ(bonds broken) - Σ(bonds formed), chemists can quickly estimate whether a proposed reaction is thermodynamically favorable (exothermic, ΔH < 0) or unfavorable (endothermic, ΔH > 0). This method is also pedagogically powerful, reinforcing the concept that chemical reactions involve rearranging atoms by breaking old bonds and making new ones.
Formula
ΔHrxn = Σ(bonds broken) - Σ(bonds formed)
Breaking bonds requires energy (+); forming bonds releases energy (-).
Common Bond Energies (kJ/mol)
C-H: 413
C-C: 348
C=C: 614
O-H: 463
C=O: 799
O=O: 498
N-H: 391
N≡N: 945
H-H: 436
Example: CH₄ + 2O₂ → CO₂ + 2H₂O
Bonds broken: 4 C-H (4×413) + 2 O=O (2×498) = 1652 + 996 = 2648 kJ
Bonds formed: 2 C=O (2×799) + 4 O-H (4×463) = 1598 + 1852 = 3450 kJ
ΔHrxn = 2648 - 3450 = -802 kJ/mol
Answer: ΔH ≈ -802 kJ/mol (exothermic)
Key Concepts
Bond Order & Bond Strength
Higher bond order = stronger bond = higher bond enthalpy. Single bonds (C-C: 348 kJ/mol) < Double bonds (C=C: 614 kJ/mol) < Triple bonds (C≡C: 839 kJ/mol).
Limitations of Average Values
Bond enthalpies are averages. For more accuracy, use standard enthalpies of formation (ΔHf°) when available. Typical errors from bond enthalpies: ±10-20 kJ/mol.
Predicting Exothermic vs Endothermic
If stronger bonds form than break, reaction is exothermic (ΔH < 0). If weaker bonds form than break, reaction is endothermic (ΔH > 0).
Common Mistakes
Forgetting to count ALL bonds
In CH₄, there are FOUR C-H bonds to break. Count bonds carefully in structural formulas.
Sign errors
Bonds broken = positive contribution; bonds formed = negative contribution. Don't flip signs.
Treating diatomic molecules as single atoms
O₂ requires breaking one O=O bond (498 kJ/mol), not two O atoms.
Notes
- Bond energies are average values; actual values vary by molecular environment.
- More accurate than bond energies: use standard enthalpies of formation.
- Useful for estimating ΔH when thermodynamic data unavailable.