Organic reactions overview

1. Organic reactions overview Dr. Clower CHEM 2411 Spring 2014 McMurry (8th 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 p 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: • Electrons move from nucleophile to electrophile • Nucleophile is negative or neutral (Nu: or Nu:-) • Electrophile is positive or neutral (E or E+) • 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 • DG/DH • DGǂ/Ea • 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 • DGǂ or Ea • 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 • DG/DH • DGǂ/Ea for each step • Intermediate

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

14. Intermediates • Which carbocation is most stable? Least stable? • Why? • Inductive effect • Donation of electrons through bonds (R groups) • 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 p bond (weaker than s bond) • HBr is a strong acid (H+ donor) with partial positive charge on H • Electrons are donated from p 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 • DG/DH • DGǂ/Ea for each step

19. Radical Reactions • Homolytic reactions • Not as common as polar reactions (heterolytic) • Mechanisms involve three steps • Initiation: start of the reaction; usually catalyzed by something • Propagation: continuation of the reaction; there can be many of these steps • 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 s bond to create reactive radicals

21. Chlorination of Methane • Propagation • Chlorine radical reacts with methane to create methyl radical • Methyl radical reacts with Cl2 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 • Radical halogenation • Combustion (alkanes as fuel) • Requires heat (Δ) or light (hn) to initiate radical formation • Chlorination (Cl2) or bromination (Br2) • 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: CH3─CH─CH2─CH3

27. Draw products for the following reactions: