The Lewis dot structures provide a picture of bonding in molecules and ions in terms of the shared pairs of electrons and the octet rule. While such a picture may not explain the bonding and behaviour of a molecule completely, it does help in understanding the formation and properties of a molecule to a large extent. Writing of Lewis dot structures of molecules is, therefor e, very useful . The Lewis dot structures can be written by adopting the following steps:
The total number of electrons required for writing the structures are obtained by adding the valence electrons of the combining atoms. For example, in the CH4 molecule there are eight valence electrons available for bonding (4 from carbon and 4 from the four hydrogen atoms).
For anions, each negative charge would mean addition of one electron. For cations, each positive charge would result in subtraction of one electron from the total number of valence electrons. For example, for the CO3 2– ion, the two negative charges indicate that there are two additional electrons than those provided by the neutral atoms. For NH4 + ion, one positive charge indicates the loss of one electron from the group of neutral atoms.
Knowing the chemical symbols of the combining atoms and having knowledge of the skeletal structure of the compound (known or guessed intelligently), it is easy to distribute the total number of electrons as bonding shared pairs between the atoms in proportion to the total bonds.
In general the least electronegative atom occupies the central position in the molecule/ion. For example in the NF3 and CO3 2–, nitrogen and carbon are the central atoms whereas fluorine and oxygen occupy the terminal positions.
After accounting for the shared pairs of electrons for single bonds, the remaining electron pairs are either utilized for multiple bonding or remain as the lone pairs. The basic requirement being that each bonded atom gets an octet of electrons.