Basic Lewis Structure "Rules"

1. Hydrogen atoms form only single bonds; they are always at the ends of a sequence of atoms. Hydrogen often is bonded to carbon, nitrogen or oxygen.
2. Oxygen atoms usually do not bond to each other, except for O₂, O₃, superoxides and peroxides.
3. Polyatomic molecules and ions often consist of a central atom (which is often listed first in the chemical formula) surrounded by terminal atoms.
4. Lewis structures tend to look symmetric.
5. To determine the total number of shared (bonding) and unshared (non-bonding) electrons in a compound use the following rule: S= N-A, where S is the total number of shared electrons, N is the total number of valence shell electrons needed by all the atoms in the molecule or ion to achieve noble gas configurations and A is the total number of electrons available in the valence shells of all the atoms in the structure. If the structure is an anion add electrons for each negative charge, if it is a cation subtract an electron for each positive charge.
6. Write the skeletal structure, and connect bonded pairs of atoms by a dash. Each dash represents a shared pair of electrons.
7. Place electrons as pairs about the outer atoms so that each (except hydrogen) has an octet.
8. Check that all atoms have satisfied the octet rule and that all the available electrons have been assigned.
9. If the central atom has fewer than 8 electrons, but all the terminal atoms are satisfied, add e^- pairs to central atom.
10. If a central atom has fewer than 8 electrons (and it is not an exception to the octet rule) a multiple bond is likely. Move electrons from outer atoms to form multiple bonds. (double is 2 shared pairs, triple is 3 shared pairs)
11. Check that your structure is the best possible arrangement of atoms using the formal charge check. (See Formal Charge Below)
12. More than one “identical structure” may have resonance.
13. Some exceptions include expanded valence (can occur in periods 3 and above), radicals (odd # of e^-) and incomplete octet (H, Be and B are common exceptions).

Formal charge is a concept used to account for the distribution of electrons in a compound. In a structure, each atom is assigned a formal charge based upon the number of valence electrons for that atom as well as the distribution of electrons in the structure. The formal charge can be calculated as,

\[FC = V - (L + {1 \over 2}S)\]

Where FC is the formal charge, V is the number of valence electrons, L is the number of lone pair electrons, S is the number of shared electrons

For more on formal charge see the "formal charge" page.

FC should be nearest to 0 for all atoms, total FC is 0 for neutral compounds, FC should = charge on polyatomic ion.

Here are a couple of videos working some examples of Lewis dot structures.