1. Chapter 3 Simple Bonding Theory Lewis Dot Structures – Resonance – Formal Charge VSEPR: the subtle effects

2. Lewis Dot Structures 1.Count valence electrons 2.Arrange atoms 3.Add bonds 4.Add lone pairs 5.Convert lone pairs to bonding pairs (octet rule and exceptions)

3. Lewis Dot Structures Examples: CO 2 SO 3 N 2 O XeF 4 ClF 3 PCl 6 –

4. Why does the octet rule work?

5. More complex NO 2 NO

6. Formal Charge = Group # - #unshared electrons on atom - # bonds to atom Example: O 3

7. Resonance Example: SO 3

8. Resonance and Formal Charge Example: SCN -

9. Resonance and Formal Charge Example: SCN -

10. Octet Rule vs. Pi Bonding Trends BeF 2 and BF 3

11. Octet Rule vs. Formal Charge Always follow octet rule Exceptions? SO 4 2-

12. VSEPR Maximize “personal space” CO 2, SO 3, SO 4 2–, PCl 5, SF 6 Lone pairs vs. bonding pairs? Single bonds vs. multiple bonds? Electronegativity effects

13. VSEPR

14. Lone Pair Effects!

15. Pi Bonds vs. Lone Pairs: Guess these bond angles.

16. Pi Bonds vs. Lone Pairs

17. Which take up more room: lone pairs or a pi bond?

18. Electronegativity Effects Molecule X-P-X Angle o PF 3 97.8 PCl 3 100.3 PBr 3 101 Explain this trend.

19. Electronegativity Effects Molecule H-X-H Angle o H 2 O 104.4 H 2 S 92.1 H 2 Se 90.6 Explain this trend. Molecule X-As-X Angle o AsF 3 AsCl 3 AsBr 3 Predict this trend.