Chemical reactions break some bonds and form others. Bond enthalpy provides an approximate energetic scale for these changes.

A bond enthalpy is the enthalpy required to break one mole of a specified bond in gaseous molecules, producing separated gaseous fragments. Tabulated values are often average bond enthalpies, because the exact energy of a bond depends on its molecular environment.

Breaking bonds requires energy, while forming bonds releases energy. This leads to the estimate

$$\boxed{\Delta H_{\mathrm{rxn}}\approx \sum D(\text{bonds broken})- \sum D(\text{bonds formed})},$$

where $D$ denotes bond enthalpy.

Example

Consider

$$\mathrm{H_2(g)+Cl_2(g)\rightarrow2HCl(g)}.$$

Using approximate average bond enthalpies

$$D(\mathrm{H-H})=436,\mathrm{kJ/mol},$$

$$D(\mathrm{Cl-Cl})=243,\mathrm{kJ/mol},$$

$$D(\mathrm{H-Cl})=431,\mathrm{kJ/mol},$$

we break one H-H bond and one Cl-Cl bond, then form two H-Cl bonds:

$$\begin{aligned} \Delta H_{\mathrm{rxn}} &\approx(436+243)-2(431)\ &=679-862\ &=-183,\mathrm{kJ}. \end{aligned}$$

The estimate is exothermic because forming the two H-Cl bonds releases more energy than is required to break the reactant bonds.

A correct bond count requires the molecular structures of reactants and products, not merely their molecular formulas.

Bond-enthalpy calculations are approximate for two important reasons. First, average bond enthalpies combine values from many molecular environments. Second, the simple bookkeeping is most naturally referenced to gaseous separated molecules, so additional intermolecular or phase-change contributions may matter when actual reactants or products are liquids, solids or solvated species.

Bond enthalpies are therefore best used to build molecular intuition and estimate reaction energetics. When accurate standard enthalpies of formation are available for the relevant physical states, those tabulated state-function data usually provide a more precise reaction enthalpy.