ions, the net charge is possessed by the ion as a whole and not by a particular atom. It is, however, feasible to assign a formal charge on each atom. The formal charge of an atom in a polyatomic molecule or ion may be defined as the difference between the number of valence electrons of that atom in an isolated or free state and the number of electrons assigned to that atom in the Lewis structure. It is expressed as : Formal charge (F.C.) on an atom in a Lewis structure = total number of valence electrons in the free atom — total number of non bonding (lone pair) electrons — ( / ) total number of bonding (shared) electrons .
. Limitations of the Octet Rule The octet rule, though useful, is not universal. It is quite useful for understanding the structures of most of the organic compounds and it applies mainly to the second period elements of the periodic table. There are three types of exceptions to the octet rule.
The incomplete octet of the central atom In some compounds, the number of electrons surrounding the central atom is less than eight. This is especially the case with elements having less than four valence electrons. Examples are LiCl, BeH and BCl . Li, Be and B have , and valence electrons only.
Some other such compounds are AlCl and BF . Odd-electron molecules In molecules with an odd number of electrons like nitric oxide, NO and nitrogen dioxide, NO , the octet rule is not satisfied for all the atoms The expanded octet Elements in and beyond the third period of the periodic table have, apart from s and p orbitals, d orbitals also available for bonding. In a number of compounds of these elements there are more than eight valence electrons around the central atom. This is termed as the expanded octet.
Obviously the octet rule does not apply in such cases. Some of the examples of such compounds are: PF , SF , H SO