and a number of coordination compounds. The counting is based on the assumption that the atom in the molecule owns one electron of each shared pair and both the electrons of a lone pair. Let us consider the ozone molecule (O ). The Lewis structure of O may be drawn as: The atoms have been numbered as , and .
The formal charge on: The central O atom marked = – – ( ) = + The end O atom marked = – – ( ) = The end O atom marked = – – ( ) = – Hence, we represent O along with the formal charges as follows: We must understand that formal charges do not indicate real charge separation within the molecule. Indicating the charges on the atoms in the Lewis structure only helps in keeping track of the valence electrons in the molecule. Formal charges help in the selection of the lowest energy structure from a number of possible Lewis structures for a given species. Generally the lowest energy structure is the one with the smallest formal charges on the atoms.
The formal charge is a factor based on a pure covalent view of bonding in which electron pairs are shared equally by neighbouring atoms. Interestingly, sulphur also forms many compounds in which the octet rule is obeyed. In sulphur dichloride, the S atom has an octet of electrons around it. Other drawbacks of the octet theory It is clear that octet rule is based upon the chemical inertness of noble gases.
However, some noble gases (for example xenon and krypton) also combine with oxygen and fluorine to form a number of compounds like XeF , KrF , XeOF etc. This theory does not account for the shape of molecules. It does not explain the relative stability of the molecules being totally silent about the energy of a molecule. .
Ionic or Electrovalent Bond From the Kössel and Lewis treatment of the formation of an ionic bond,