1. Formal Charge Rules
Often, many Lewis dot structures are possible for the same compound. These are called resonance structures, but often we should write a reasonable one, which is stable. The formal charge guides us about the stability of the dot structure. The guidance are called formal charge rules:
Formulas with the lowest magnitude of formal charges are more stable.
More electonegative atoms should have negative formal charges.
Adjacent atoms should have opposite formal charges.

2. Let’s Draw Lewis Dot Structures
Acetate ion, CH3CO2–
   
   
Required Valence Electrons:
For Atoms
Charge (-1)
Total 23
1 EXTRA 24
Electrons As Written:
Single Bonds
Lone Pair
Double/Triple Bonds
Still Needed 12

3. Let’s Draw Lewis Dot Structures.. -1 .. .. .. .. -1 .. .. .. .. ..
Required Valence Electrons:
For Atoms
Charge (-1)
Total 23
1 EXTRA 24
Electrons As Written:
Single Bonds
Lone Pair
Double/Triple Bonds
Still Needed 12 10 1
The two carbon-oxygen bonds have equal length and equal strength; they are identical! We have two resonance structures.

4. Let’s Draw Lewis Dot Structures
The classic example of a resonance structure: O3.
Both structures are correct!

5. Let’s Draw Lewis Dot Structures
In this example, O3 has two resonance structures:
Conceptually, we think of the bonding being an average of these two structures.
Electrons are delocalized between the oxygens such that on average the bond strength is equivalent to 1.5 O-O bonds.

6. Structural Isomers
What if different sets of atomic linkages can be used to construct correct Lewis dot structures
Both are correct, but which is likely to be more stable?

7. Formal Charge: Compare the nuclear charge (+Z) to the number of electrons (dividing bonding electron pairs by 2). Difference is known as the “formal charge”. #e Z
Formal Charge
Structure with less formal charge is more stable.

8. Example: CO2 e Z FC
More Stable

9. Beyond the Octet Rule Numerous exceptions to the octet rule occur!
Let me mention three classes of violations:
Sub-octet systems (electron deficient compounds)
Valence shell expansion (hypervalency)
Odd-electron systems (radicals)

10. Beyond the Octet Rule (cont.)
Some atoms (Be and B in particular) undergo bonding, but will form stable molecules that do not fulfill the octet rule.
Experiments demonstrate that the B-F bond strength is consistent with single bonds only.

11. Beyond the Octet Rule (cont.)
For third-row elements (“Period 3”), the energetic proximity of the d orbitals allows for the participation of these orbitals in bonding.
When this occurs, more than 8 electrons can surround a third-row element.
Example: ClF3 (a 28 e- system)
IF7 is the winner!
F obey octet rule
Cl has 10 e-

12. Beyond the Octet Rule (cont.)
Finally, one can encounter odd-electron systems (radicals) where full pairs cannot exist.
Example: Chlorine Dioxide.
Unpaired electron
Most radicals are chemically unstable, that is,
highly reactive. They pair up with themselves as
well as react with themselves and other
substances.

13. De re

15. Harpoon Model (Herschbach)
The reaction of M, an alkali atom, with XY a halogen molecule:
M + XY  MX + Y
has a relatively large cross section (meaning that the reaction takes place very efficiently). WHY???

16. The valence electron of the electropositive atom jumps at a relatively large distance to the electronegative molecule.
The resulting attractive Coulomb force creates an ionic bond. Because the negative molecular ion is vibrationally excited, it either immediately dissociates, or is easily dissociated by the Coulomb force, so that M+X- can be formed. The metaphor of the harpoon model is that of the small alkali atom (Captain Ahab) catching a big halogen atom (Moby Dick), by throwing its valence electron serving as a harpoon and hauling it in with the Coulomb force of attraction.
M + XY  M+ + (XY)-