Drawing Lewis Structures

Procedure

1.) Count up the total number of valence electrons available. Add electrons for negatively charged species and subtract electrons for positively charged species.

Ex: CH₂O (C is central atom)

4 e^-(C) + 1 e^-(H) + 1 e^-(H) + 6 e^-(O) = 12 e^-

2.) Calculate the total number of electrons that would be needed if each atom had its own noble-gas shell of electrons (two for hydrogen, eight for all else).

Ex: CH₂O

8 e^-(C) + 2 e^-(H) + 2 e^-(H) + 8 e^-(O) = 20 e^-

3.) Calculate the number of bonding electrons by subtracting the number in step 1 from the number in step 2.

Ex: CH₂O

20 e^- - 12 e^- = 8 e-

4.) Assign two bonding electrons (one pair) to each bond.

Ex: CH₂O

5.) Assign the remaining bonding electrons (in pairs) by making double and triple bonds. In general, C, N, O, S form double bonds and C, N form triple bonds.

Ex: CH₂O

used 6 e^- in step 4, so there are 2 bonding e^- remaining

6.) Assign the remaining electrons to lone pairs, giving each atom (except for hydrogen) an octet.

Ex: CH₂O

7.) Determine the formal charge of each atom and write any non-zero charges next to the atom. The formal charges must add up to the overall charge on the molecule.

Ex: CH₂O

C: 4 - 4 = 0

H: 1 - 1 = 0 (same for both)

O: 6 - 4 - 2 = 0

total = 0 + 0 + 0 = 0; molecule is neutral

Exceptions:

1.) Odd Electron Molecules:

Any molecule with an odd number of electrons cannot satisfy the octet rule.

Ex: CH₃

2.) Octet Deficient Molecules:

Some molecules are stable even though not all their atoms have a full octet.

Ex: AlCl₃

3.) Valence Shell Expansion:

Atoms below the second row can expand their valence shells and accept more than eight electrons, if needed.

Ex: SO₃

Resonance:

1.) Occurs whenever there are two or more equally good Lewis structures for a molecule.

2.) The actual structure of the molecule is an average of the Lewis structures.

Example: SO₃^2-