The Bond Enthalpy is the energy required to break a chemical bond. It is usually expressed in units of kJ mol-1, measured at 298 K. The exact bond enthalpy of a particular chemical bond depends upon the molecular environment in which the bond exists. Therefore, bond enthalpy values given in chemical data books are averaged values.
Bond breaking is an endothermic process, and the bond enthalpy involved is given a +ve sign. Bond making is an exothermic process. For the formation of a given chemical bond, the bond enthalpy has the same value, except it is given a -ve sign.
Generally, the shorter a chemical bond the stronger it is. This can be illustrated with the halogen-halogen bond enthalpies:
| Bond | Bond Length (nm) | Bond Enthalpy (kJ mol-1)
| F – F | 0.142 | 158
| Cl – Cl | 0.199 | 242
| Br – Br | 0.228 | 193
| I – I | 0.267 | 151
| |
Note the apparently anomolous bond enthalpy of the F-F bond. This is explained by the lone pairs of electrons on the fluorine atoms being closer together and therefore repelling each other more than is the case in the other halogen molecules.
Bond enthalpies can be used to calculate enthalpy changes for reactions, but even though standard bond enthalpy values are used, the calculated DH is likely to differ from that given in a chemical data book. There are two reasons for this. Bond enthalpy values are averaged values, and in the chemical reactions where they are applied, all the reactants and products are taken to be gaseous.
In photochemical reactions, for example, the free radical chlorination of methane, a chlorine molecule, Cl2, absorbs a photon of ultra-violet radiation sufficient in energy to break the Cl-Cl bond homolytically. Given the bond enthalpy of the Cl-Cl bond, E(Cl-Cl) = 242 kJ mol-1, the energy required to break just one such bond can be calculated.
Energy to break one Cl-Cl bond = 242 x 1000/6.023 x 1023 = 4.02 x 10-19 J