reason for some difference in energy of the same O – H bond in different molecules like C H OH (ethanol) and water. Therefore in polyatomic molecules the term mean or average bond enthalpy is used. It is obtained by dividing total bond dissociation enthalpy by the number of bonds broken as explained below in case of water molecule, Average bond enthalpy = + = . kJ mol – .
. Bond Order In the Lewis description of covalent bond, the Bond Order is given by the number of bonds between the two atoms in a molecule. The bond order, for example in H (with a single shared electron pair), in O (with two shared electron pairs) and in N (with three shared electron pairs) is , , respectively. Similarly in CO (three shared electron pairs between C and O) the bond order is .
For N , bond order is and its is kJ mol – ; being one of the highest for a diatomic molecule. Isoelectronic molecules and ions have identical bond orders; for example, F and O – have bond order . N , CO and NO + have bond order . A general correlation useful for understanding the stablities of molecules is that: with increase in bond order, bond enthalpy increases and bond length decreases.
. . Resonance Structures It is often observed that a single Lewis structure is inadequate for the representation of a molecule in conformity with its experimentally determined parameters. For example, the ozone, O molecule can be equally represented by the structures I and II shown below: In both structures we have a O–O single bond and a O=O double bond.
The normal O–O and O=O bond lengths are pm and pm respectively. Experimentally determined oxygen-oxygen bond lengths in the O molecule are same ( pm). Thus the oxygen-oxygen bonds in the O molecule are intermediate between a double and a single bond. Obviously, this cannot be represented by either of the two Lewis structures shown above.