1. ORGANIC REACTIONS OVERVIEW Dr. Clower CHEM 2411 Spring 2014 McMurry (8 th ed.) sections 6.1, 6.2, 6.4-6, 6.8-10, 7.10, 10.8

2. Organic Reactions Types of Reactions: Addition Elimination Substitution Rearrangement Oxidation Reduction See handout

3. Reaction Mechanisms The details of how reactions occur Bonds broken Bonds formed Electron rearrangement Order of steps Kinetics (rate) Thermodynamics (energy) Role of solvent, catalysts, etc.

4. Bond Breaking Symmetrical/radical/homolytic One electron to each atom Fishhook arrow Result in formation of free radicals Unsymmetrical/polar/heterolytic Both electrons to one atom Regular curved arrow Electrons move to more electronegative atom

5. Bond Formation Symmetrical/radical/homogenic One electron from each atom Unsymmetrical/polar/heterogenic Both electrons from one atom What is the nucleophile? What is the electrophile?

6. Nucleophiles and Electrophiles Nucleophile Electron pair donor Contain lone pair or  e - Electrophile Electron pair acceptor Positive or partial positive charge Remember electrons always move from nucleophile to electrophile

7. Nucleophile or Electrophile?

8. Drawing Mechanisms Curved arrows Some guidelines: 1. Electrons move from nucleophile to electrophile 2. Nucleophile is negative or neutral (Nu: or Nu: - ) 3. Electrophile is positive or neutral (E or E + ) 4. Obey the octet rule See Mechanisms worksheet

9. Energy Diagrams Change in energy as reaction proceeds A one-step reaction: Label: Axes Starting material Product Transition state  G/  H  G ǂ /E a Where does bond breaking occur? Where does bond making occur? How do you know if the reaction is endothermic or exothermic?

10. Transition State One transition state per step Highest energy species in the step Unstable; cannot be isolated Resembles species (starting material or product) that is closest in energy Hammond’s postulate In an endothermic step the TS resembles the product In an exothermic step the TS resembles the reactant/starting material

11. Activation Energy  G ǂ or E a Energy difference between starting material and transition state Minimum energy needed for reation to occur High activation energy = slow reaction Rate-determining step (RDS) The slowest step The step with the largest activation energy

12. Energy Diagrams A two-step reaction: Label: Axes Starting material Product Transition states  G/  H  G ǂ /E a for each step Intermediate

13. Energy minimum between two transition states Higher energy than starting material or product Usually cannot isolate (unstable) Types of intermediates: 1. Free radicals 2. Carbocations

14. Intermediates Which carbocation is most stable? Least stable? Why? 1. Inductive effect Donation of electrons through bonds (R groups) 2. Hyperconjugation Donation of electrons through orbitals Other stable carbocations are resonance-stabilized

15. An Example Reaction HBr + ethylene → bromoethane What type of reaction is this? What is the nucleophile? Electrophile? Look at structure: Ethylene C=C has high electron density (4 e - ); relatively easy to break  bond (weaker than  bond) HBr is a strong acid (H + donor) with partial positive charge on H Electrons are donated from  bond of ethylene to H of HBr Sigma bond of ethylene is not broken

16. Mechanism Two steps Step 1: Step 2:

17. Mechanism The mechanism can be written as one scheme:

18. Energy Diagram Label: Axes Starting material Intermediate Product  G/  H  G ǂ /E a for each step

19. Radical Reactions Homolytic reactions Not as common as polar reactions (heterolytic) Mechanisms involve three steps 1. Initiation: start of the reaction; usually catalyzed by something 2. Propagation: continuation of the reaction; there can be many of these steps 3. Termination: end of the reaction An example reaction: chlorination of methane What type of reaction is this?

20. Chlorination of Methane Initiation Caused by irradiation with UV light Break  bond to create reactive radicals

21. Chlorination of Methane Propagation Chlorine radical reacts with methane to create methyl radical Methyl radical reacts with Cl 2 to give product and more Cl radical New Cl radical repeats this propagation process (a chain reaction)

22. Chlorination of Methane Termination Two radicals collide to form stable product Break the reaction cycle

23. Radical Halogenation Used to synthesize alkyl halides from alkanes One of only two alkane/cycloalkane reactions 1. Radical halogenation 2. Combustion (alkanes as fuel) Requires heat (Δ) or light (h ) to initiate radical formation Chlorination (Cl 2 ) or bromination (Br 2 ) Iodine is too endothermic; fluorine is too reactive Typically results in mixtures of products

24. Halogenation of Alkanes Ex: ethane Ex: butane Why is this? Consider the intermediate structure…

25. Halogenation of Alkanes Substitution is favored at more substituted carbons Tertiary > secondary > primary The tertiary radical is more stable than the secondary radical Regiochemistry

26. Stereochemistry of Halogenation If the product contains a stereocenter, what is the stereochemistry? This reaction will produce a racemic mixture. Why? Look at radical intermediate: CH 3 ─CH─CH 2 ─CH 3

27. Draw products for the following reactions: a) b) c)