RULES FOR LEWIS STRUCTURES
A Lewis structure consists of the electron distribution in a compound and the formal charge on each atom. You are expected to be able to draw such structures to represent the electronic structure of compounds. The following rules are given to assist you.
(FORMAL CHARGE IS NOT REQUIRED IN YOUR CAMBRIDGE AS EXAM)
1. Determine whether the compound is covalent or ionic. If covalent, treat the entire molecule. If ionic, treat each ion separately. Compounds of low electronegativity metals with high electronegativity nonmetals (ΔEN > 1.6) are ionic as are compounds of metals with polyatomic anions. For a monoatomic ion, the electronic configuration of the ion represents the correct Lewis structure. For compounds containing complex ions, you must learn to recognize the formulas of cations and anions.
2. Determine the total number of valence electrons available to the molecule or ion by:
(a) summing the valence electrons of all the atoms in the unit and
(b) adding one electron for each net negative charge or subtracting one electron for each net positive charge. Then divide the total number of available electrons by 2 to obtain the number of electron pairs (E.P.) available.
3. Organize the atoms so there is a central atom (usually the least electronegative) surrounded by ligand (outer) atoms. Hydrogen is never the central atom.
4. Determine a provisional electron distribution by arranging the electron pairs (E.P.) in the following manner until all available pairs have been distributed:
a) One pair between the central atom and each ligand atom.
b) Three more pairs on each outer atom (except hydrogen, which has no additional pairs), yielding 4 E.P. (i.e., an octet) around each ligand atom when the bonding pair is included in the count.
c) Remaining electron pairs (if any) on the central atom.
5. Calculate the formal charge (F) on the central atom.
a) Count the electrons shared as bonds. Total = b
b) Count the electrons owned as lone pairs. Total = n
c) F = V – (n + b/2), where V = number of valence electrons for the atom.
6. If the central atom formal charge is zero or is equal to the charge on the species, the provisional electron distribution from (4) is correct. Calculate the formal charge of the ligand atoms to complete the Lewis structure.
7. If the structure is not correct, calculate the formal charge on each of the ligand atoms. Then to obtain the correct structure, form a multiple bond by sharing an electron pair from the ligand atom that has the most negative formal charge.
a) For a central atom from the second (n = 2) row of the periodic table continue this process sequentially until the central atom has 4 E.P. (an octet).
b) For all other elements, continue this process sequentially until the formal charge on the central atom is reduced to zero or two double bonds are formed.
8. Recalculate the formal charge of each atom to complete the Lewis structure.
Written by Patrick A. Wegner; California State University, Fullerton.