Bond Enthalpies and Estimating Enthalpy Change

A bond exists because the bonded arrangement is lower in energy. Pulling the atoms apart requires that energy back, so bond dissociation is always endothermic and the tabulated values are always positive. Forming the same bond releases the same magnitude.

Separating bonded gaseous atoms raises their energy; a positive bond-dissociation enthalpy is the cost of that specified process. Recombining those same atoms to restore the same bond has the opposite sign. Burning a fuel releases energy from the net replacement of bonds, not from breaking the fuel's bonds alone. Melting ice also should not be described as breaking its covalent O–H bonds: that phase change chiefly alters intermolecular interactions.

Every reactant bond that changes must be broken and every product bond must be formed. Adding the costs and subtracting the payback gives the net enthalpy change.

Imagine taking gaseous reactants to separated atoms, then assembling gaseous products. The first leg adds bond-breaking energies, the second subtracts the listed dissociation energies for product bonds. That Hess cycle explains ΔH ≈ ΣD(broken) − ΣD(formed). It is not the real reaction mechanism and does not imply every bond actually breaks simultaneously. Each term must include both the number of such bonds per molecule and the molecule's equation coefficient.

Draw or visualise the structures. Methane has four C-H bonds, not one. O₂ is a double bond, N₂ is a triple bond, and CO₂ contains two C=O bonds. Miscounting the structures, rather than the arithmetic, is the usual source of error.

For H₂ + Cl₂ → 2 HCl, the ledger contains one H–H and one Cl–Cl broken, two H–Cl formed. With 436, 242 and 431 kJ/mol, ΔH ≈ 436 + 242 − 2(431) = −184 kJ. For methane combustion to gaseous water, count four C–H and two O=O broken; two C=O and four O–H formed. A coefficient of two waters means four O–H bonds, not two. Triple and double bonds use their own tabulated D values rather than three or two copies of a single-bond value.

A C-H bond in methane is not identical to a C-H bond in ethanol, so tables list an average across many compounds. The method also assumes every species is gaseous, because condensing to a liquid releases additional energy that bond enthalpies do not capture.

Average bond values vary with the surrounding molecule, so cancelling “unchanged C–H bonds” is an approximation when their environments change. For ethene hydrogenation, the useful net ledger replaces C=C and H–H with C–C and two C–H bonds; it estimates the gas-phase heat. To estimate a liquid product, add the appropriate condensation or other phase correction. Formation enthalpies for the exact molecules and states generally provide the better thermochemical result when available. A difference between two methods can arise from both averaging and phase choice; do not assign all discrepancy to one cause without checking.