Drawing Lewis Structures
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).
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.
20 e- - 12 e- = 8 e-
4.) Assign two bonding electrons (one pair) to each bond.
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.
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.
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.
C: 4 - 4 = 0
H: 1 - 1 = 0 (same for both)
O: 6 - 4 - 2 = 0
total = 0 + 0 + 0 = 0; molecule is neutral
1.) Odd Electron Molecules:
Any molecule with an odd number of electrons cannot satisfy the octet rule.
2.) Octet Deficient Molecules:
Some molecules are stable even though not all their atoms have a full octet.
3.) Valence Shell Expansion:
Atoms below the second row can expand their valence shells and accept more than eight electrons, if needed.
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.