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:  SO32-