bond-dissociation energy

The bond dissociation energy, \[D_{0}\], is defined as the standard enthalpy change when \[\ce{R-X}\] is cleaved by homolysis or heterolysis to give fragments \[\ce{R}\] and \[\ce{X}\]. Officially, the IUPAC definition of bond-dissociation energy refers to the energy change that occurs at 0 K. However, bond dissociation energy is generally used, albeit imprecisely, interchangeably with the bond dissociation enthalpy, \[DH^{\circ}_{298}\], which generally refers to the enthalpy change at room temperature (298 K). Although there are technically differences between bond dissociation energies at 0 K and 298 K, those difference are not large and generally do not affect interpretations of chemical processes.

It is derived via the thermochemical equation: \[D_{0}(\ce{R-X})=\Delta H^{\circ}_{f}(\ce{R})+\Delta H^{\circ}_{f}(\ce{X})-\Delta H^{\circ}_{f}(\ce{RX})\], where \[D_{0}(\ce{R-X})\] is the bond disassociation energy, and \[\Delta H^{\circ}_{f}\] is the enthalpy of formation. This equation tells us that the bond dissociation energy for a given bond is equal to the energy of the individual components that make up the bond when they are free and unbonded minus the energy of the components when they are bonded together. Essentially, the bond dissociation energy describes the energy required to break said bond (which is equal to the energy released when the bond was formed).

As a typical example, the bond-dissociation energy for one of the \[\ce{C-H}\] bonds in ethane is defined as the standard enthalpy change of the process:
\[\ce{CH3CH2-H -> CH3CH2^{.} + H^{.}}\], \[DH^{\circ}_{298}(\ce{CH3CH2-H})=423.0\pm 1.7\text{kJmol$^{-1}$}\]