Covalent chemical bonds represent a repository of potential energy. This bond energy (chemists tend to refer to it as bond dissociation energy) defines the strength of the covalent bond. The stronger the bond, the greater the bond energy. Bond energy is defined as the change in enthalpy (H) when one mole of chemical bonds is broken in the gaseous state. The equation below describes the bond energy for one mole of single covalent bonds.

For example:

HF_(g) -> H_(g) + F_(g)    ΔH = +565 kJ

From this data we can deduce that ΔH for the H-F bond is 565 kJ.
Note: The stoichiometry of this equation has been adjusted to emphasize 1 mole of H-F bonds.

Determining the energy change for a chemical reaction (enthalpy change) implies subtracting the energy released when bonds are made in products from the energy required to break the bonds of the reactants.

Ionic and other bond energies can be discussed in terms of energy diagrams that relate energy to distance of separation.

Bond energies are not fixed, but vary somewhat from case to case. Nevertheless, tables of average bond enthalpies are available.