1. Next Steps: Resonance

2. Formal Charge When atoms do not exhibit ‘normal’ bonding patterns, they will contain a ‘formal charge’. Formal Charge does not indicate an actual ionic charge – it indicates the distribution of electrons

3. Dimethyl Sulfoxide (DMSO) Normally, Sulfur owns 6 valence electrons, but in this structure, it only owns 5 Therefore, Sulfur has formally lost 1 electron and has a + charge Likewise, Oxygen normally owns 6 valence electrons – in this structure it owns 7, so it has a formal - charge

4. Calculating Formal Charge FC = #valence e - - [(1/2 bonded e - ) + nonbonding e - ] Easier Calculation: FC = #valence e - - bonds – dots You Try It: Calculate any fc’s for nonhydrogen atoms H 3 C-C ≡ N-O

5. Note: From now on, lone pairs or formal charges must be shown when needed. You may show both, but it is not necessary. Atoms that exhibit normal bonding patterns may assumed to have a formal charge of zero Read pages 10-19 & try problems

6. Resonance This is why we study formal charge: Consider Nitromethane:

7. Nitromethane EPM Experiments show that each N-O bond is equivalent. Examine electron distribution:

8. Why? The true structure is a resonance hybrid. The electrons are distributed evenly with both oxygen atoms bearing equal negative charge. Remember: ◦ Resonance structures are not real. They only help us to envision electron distribution. Only by knowing the contributing structures can we envision the real structure.

9. Benzene

10. 2 Major Rules for Resonance 1. Never break a single bond 2. Never exceed an octet for 2 nd row elements For more practice see handout problems 2.2 – 2.12 pgs 26- 27

11. Drawing Arrows to Show Movement of Electrons: Pushing Electrons Where the electrons come from Where the electrons are moving to

12. Example:

13. You try it Draw arrows that show how one structure becomes the other through resonance: More problems: pg 29; 2.14 – 2.19

14. Patterns for Drawing Resonance Structures: 1. Lone pair next to pi bond 2. Lone pair next to a positive charge 3. Pi bond next to a positive charge 4. Pi bond between two atom where one of those is electronegative 5. Pi bonds going all the way around a ring 6. Pi bond next to a free radical

15. 1. Lone Pair Next to a Pi Bond “Next to” – a lone pair is separated from a pi bond by exactly one single bond

16. 2. Lone pair next to + charge Remember a + charge means that there is less electron density than usual, so there is an empty orbital available

17. Example:

18. 3. Pi Bond next to + charge +

19. 4. Pi Bond between two atoms where one is electronegative An electronegative atom can support an additional pair of electrons and a formal negative charge

20. 5. Pi bonds going all the way around a ring

21. Phenanthrene How many resonance structures for this example?

22. 6. Pi bond next to free radical What is a free radical? o radical - (free radical) a neutral substance that contains a single, unpaired electron in one of its orbitals, denoted by a dot (·) leaving it with an odd number of electrons. o Radicals are highly reactive and unstable o Radicals can form from stable molecules and can also react with each other.

23. Showing resonance of free radicals Use half-arrows to represent the movement of single electrons

24. You try it Show all of the resonance forms for the following structure:

25. A look at Pyridine The lone pair on the nitrogen does not participate in resonance due to its position in an sp 2 hybrid orbital

26. Draw All Resonance Structures for Pyridine

27. Significant Resonance Structures Not all resonance structures are significant. Three rules help us choose structures that are significant ◦ 1. Minimize Charges ◦ 2. Electronegative atoms can bear positive charge only if they have a full octet ◦ 3. Avoid resonance structures in which two carbon atoms bear opposite charges

28. 1. Minimize Charges

29. 2. Electronegative Atoms & positive charge

30. 3. Avoid Carbons with opposite charges