1. 1 Organic Chemistry, Third Edition Janice Gorzynski Smith University of Hawai’i Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Chapter 6 Lecture Outline Prepared by Layne A. Morsch The University of Illinois - Springfield

2. 2 1. Organic Chemistry. - Structure & Reactions 2. Chemical reactions: 1) Bond break & bond formation 2) Nu: - (electron-rich) reacts with E + (electron- deficient) 3) Learn by a few basic theme (principle). 3. Reaction Mechanism: How bonds are broken and formed as starting material is converted into product. Organic Reactions

3. 3 Ways to Write Organic Reactions The symbols “h ” and “  ” are used for reactions that require light or heat, respectively. Reactant  Product.

4. 4 -The first step occurs before the second step. -The reagents are added in sequence. Writing Equations for Sequential Reactions

5. Kinds of Organic Reactions 1.Substitution 2.Elimination 3.Addition 5

6. 6 Substitution Reactions One  bond breaks and another forms at the same carbon atom.

7. 7 Two  bonds are broken, and a  bond is formed. Elimination Reactions

8. 8 A  bond is broken and two  bonds are formed. Addition Reactions

9. 9 Addition and elimination reactions are exactly opposite. Often these reactions are reversible. Relationship of Addition and Elimination Reactions

10. 10 Reaction mechanism: How bonds are broken and formed as starting material is converted into product. A reaction can occur either in one step or a series of steps. Reaction Mechanisms

11. 11 Only two ways to break (cleave) a bond: - Homolysis & Heterolysis. Bond Breaking – Homolysis & Heterolysis

12. 12 1. Bond Breaking – Homolysis Use a half-headed curved arrow (fishhook). Homolysis generates highly unstable radicals

13. 13 2. Bond Breaking – Heterolysis The more electronegative atom takes 2e -. Heterolysis generates a carbocation or a carbanion. Carbocation & carbanion are unstable intermediates  reactive

14. 14 Figure 6.2 Reactive Intermediates Resulting from Breaking a C-Z Bond

15. 15 Radicals and carbocations: electrophile Carbanions: nucleophile Bond Breaking – Intermediates

16. 16 1. 2 different ways. 1) radical + radical 2) Nu: - + E + 2. Bond formation always releases energy. Bond Forming

17. 17 Arrows Used in Organic Reactions

18. 18 Bond Dissociation Energy Homolysis: endothermic

19. 19

20. 20 Stronger the bond  higher bond dissociation energy. Generally, shorter bonds are stronger bonds (= stable). Bond Dissociation Energy and Bond Strength

21. 21 1) Bond dissociation energies are determined in the gas phase 2) Most organic reactions occur in a liquid solvent  Solvation energy contributes to ∆H 0 overall of a reaction. 3) Bond dissociation energy  a useful approximation to calculate energy changes in a reaction. Limitations on Bond Dissociation Energies

22. 22 Bond dissociation energies are used to calculate the enthalpy change (  H ° ) in a reaction. Enthalpy Change in Reactions ∆H 0 overall < 0: stronger bond formation (lower energy) ∆H 0 overall > 0: weaker bond formation (higher energy)

23. 23 Determining  H ° for a Reaction Consider bonds being broken and formed.

24. 24  H° is negative for both oxidations (exothermic).  Bonds in the products are stronger than bonds in the reactants. Enthalpy Changes in Oxidation Reactions

25. 25 For a reaction to be practical, thermodynamics and kinetics should be considered. 1)Thermodynamics: the relative amounts of reactants and products at equilibrium. 2) Kinetics: reaction rate (how fast a reaction occurs). Thermodynamics and Kinetics

26. 26 The equilibrium constant, K eq Thermodynamics & K eq K eq > 1: spontaneous rx. (favor product formation) K eq < 1: non-spontaneous rx.

27. 27  G° = -RT ln K eq K eq > 1   G° < 1 : product formation. K eq 1 : equilibrium favors the reactants. K eq & Free Energy (  G°)

28. 28 Figure 6.3 K eq & Free Energy (  G°) and Free Energy  G° is the overall energy difference between reactants and products.

29. 29  G° = - RT ln K eq K eq &  G°

30. 30 Knowing the energy difference between two conformations permits the calculation of the amount of each at equilibrium. Conformations and Equilibrium

31. 31 ∆G 0 : free energy change, total energy change ∆H 0 : enthalpy, change in bonding energy ∆S 0 : entropy, change in disorder. 1) In most reactions, T  S° <<  H°), Total Energy Change (  G°) 2) But, T  S° >> 0, ∆G 0 = ∆H 0 + T∆S 0

32. 32 Entropy (  S° ): a measure of the randomness of a system.  S° > 0 : more disordered  spontaneous rx.  S° are important when: 1) Changes in # of mol. in the chemical equation 2) Acyclic mol.  cyclic mol.  S° in Reactions

33. 33 1.Energy changes as reactants are converted to products. 2.Transition state - a partial bond - an unstable energy maximum. 3. Energy of activation, E a : the energy difference between the transition state and the reactant. Energy Diagrams

34. 34 For the general reaction: The energy diagram would be shown as: Energy Diagrams rate, kinetics equilibrium, thermodynamics

35. 35 Kinetics: reaction rates. E a is the energy barrier. Kinetics and Energy Diagrams - E a, concentration, temperature  rate -  G °,  H °, and K eq  equilibrium

36. 36 Comparing ∆H° and E a in Energy Diagrams

37. 37 1. Larger E a : more energy to break bonds  slower rate. 2. Transition state: - partial bond - Has a partial charge (δ -, δ + ) - Designation: [ ] ‡, double dagger ( ‡ ). Transition States

38. 38 Slow, Endothermic Energy Diagram

39. 39 Slow, Exothermic Energy Diagram

40. 40 Fast, Endothermic Energy Diagram

41. 41 Fast, Exothermic Energy Diagram

42. 42 Each step has energy diagram. Energy Diagrams and Two-Step Reactions

43. 43 Stepwise Reaction Energy Diagram–Step 1

44. 44 Stepwise Reaction Energy Diagram–Step 2

45. 45 Figure 6.6 Stepwise Reaction Overall Energy Diagram

46. 46 Kinetics: reaction rates. E a : Activation energy (energy barrier). Kinetics and Energy Diagrams - E a, concentration, the temperature  rate -  G °,  H °, and K eq  equilibrium

47. 47 1. Rate equation 2. Rate constant (k) 3. experimentally determined - large rate constant  fast reaction - small rate constant  slow reaction Rate Equation

48. 48 v = k [reactants] K= ƙ T/ ɦ e -∆G ǂ /RT v = k[S] = [S] ƙ T/ ɦ e -∆G ǂ /RT ƙ : Boltzman constant, T : Plank constant 1)higher E a (∆G ǂ )  smaller k. 2)2) Higher conc.  faster rate (v). 3) The order of reactions  related reaction mechanism (1) 2 nd order rx.: one-step rx., bimolecular rx. (2) 1 st order rx.: step-wise rx., unimolecular rx. Rate equations

49. 49 Rate-determining step in a multistep reaction. Only the rate-determining step appears in the rate equation. Rate Equation for Two-Step Reaction

50. 50 Catalyst: accelerate the rate of a reaction. 1) lower E a 2) the position of equillibium is unaffected. Figure 6.7 Catalyst

51. 51 Enzyme: Biological Catalyst ex) Lactase