An Overview of Organic Reactions

1. An Overview of Organic Reactions Why and how chemical reactions take place? How a reaction can be described?

2. Kinds of Organic Reactions Addition reactions – two molecules combine

3. Elimination reactions – one molecule splits into two

4. Substitution – parts from two molecules exchange

5. Rearrangement reactions – a molecule undergoes changesin the way its atoms are connected

6. How Organic Reactions Occur: Mechanisms • In a clock the hands move but the mechanism behind the face is what causes the movement • In an organic reaction, we see the transformation that has occurred. Themechanism describes the steps behind the changes that we can observe • Reactions occur in defined steps that lead from reactant to product

7. Steps in Mechanisms • We classify the types of steps in a sequence • A step involves either the formation or breaking of a covalent bond • Steps can occur individually or in combination with other steps • When several steps occur at the same time they are said to be concerted

8. Types of Steps in Reaction Mechanisms Bond formation or breaking can be symmetrical or unsymetrical Symmetrical-homolytic Unsymmetrical-heterolytic

9. Indicating Steps in Mechanisms • Curved arrows indicate breaking and forming of bonds • Arrowheads with a “half” head (“fish-hook”) indicate homolytic and homogenic steps (called ‘radical processes’) • Arrowheads with a complete head indicate heterolytic and heterogenic steps (called ‘polar processes’)

10. Radical Reactions A radical can break a bond in another moleculeandabstract a partnerwith an electron, giving substitution in the original molecule

11. Steps in Radical Substitution Initiation – homolytic formation of two reactive species with unpaired electrons Propagation – reaction with molecule to generate radical Termination – combination of two radicals to form a stable product:

12. Polar Reactions • Molecules can contain local unsymmetrical electron distributions due to differences in electronegativities • This causes a partial negative charge on an atom and a compensating partial positive charge on an adjacent atom • The more electronegative atom has the greater electron density • Elements such as O, F, N, Cl are more electronegative than carbon

14. Polarity patternsin some functional groups

15. Polar reactions occur between regions of high electron density and regions of low electron density Polarizability is the tendency to undergo polarization

16. Polar Reactions An electrophile, an electron-poor species, combines with a nucleophile, an electron-rich species An electrophile is a Lewis acid A nucleophile is a Lewis base

17. Some nucleophiles and electrophiles

18. An Example of a Polar Reaction: Addition of HBr to Ethylene • HBr adds to the  part of C-C double bond • The  bond is electron-rich, allowing it to function as a nucleophile • H-Br is electron deficient at the H since Br is much more electronegative, making HBr an electrophile

19. Addition of HBr to Ethylene - Mechanism

20. Using Curved Arrows in Polar Reaction Mechanisms • Curved arrows are a way to keep track of changes in bonding in polar reaction • The arrows track “electron movement” • Electrons always move in pairs • Charges change during the reaction • One curved arrow corresponds to one step in a reaction mechanism • The arrow goes from the nucleophilic reaction site to the electrophilic reaction site

21. The nucleophilic site can be neutral or negatively charged

22. The electrophilic site can be neutral or positively charged

23. Describing a Reaction: Equilibria Reactions can go either forward or backward to reach equilibrium • The multiplied concentrations of the products divided by the multiplied concentrations of the reactant is the equilibrium constant, Keq • Each concentration is raised to the power of its coefficient in the balanced equation.

24. If the value of Keq> 1, this indicates that at equilibrium most of the material is present as products If the value of Keq< 1, this indicates that at equilibrium most of the material is present as substrates WhenKeq> 1000, reaction is considered „complete” – amount of reactant left less than 0.1%

25. Equilibrium and Free Energy • The ratio of products to reactants is controlled by their relative Gibbs free energy • This energy is released on the favored side of an equilibrium reaction • The change in Gibbs free energy between products and reactants is written as “ΔG” • If Keq > 1, energy is released to the surroundings (exergonic reaction) • If Keq < 1, energy is absorbed from the surroundings (endergonic reaction)

26. Relationship of Keq and Free Energy Change • The standard free energy change at 1 atm pressure and 298 K is ΔGº • The relationship between free energy change and an equilibrium constant is: • ΔGº = - RT ln Keq • where • R = 1.987 cal/(K x mol) gas constant • T = temperature in Kelvin • ln Keq = natural logarithm of Keq

27. Describing a Reaction: Thermodynamic parameters • The Gibbs free energy change is attributable to a combination of two factors – an enthalpy factor ΔHo and an entropy factor ΔSo • ΔGº = ΔHº - TΔSo • ΔHº - enthalpy of reaction (change in total bonding energy during a reaction) • ΔSo- entropy of reaction (change in in the amount of molecular disorder caused by reaction)

29. Bond Dissociation Energies Bond dissociation energy (D): amount of energy required to break a givenbondto produce two radical fragments when the molecule is in the gas phase at 25˚ C Changes in bonds can be used to calculate net changes in heat (Enthalpy = ΔH)

30. Bond Dissociation Energies

31. Bond Dissociation Energies

32. Addition of HBr to Ethylene - Mechanism

33. Energy Diagrams for Single-step Reaction • The highest energy point in a reaction step is called the transition state • The energy needed to go from reactant to transition state is the activation energy (ΔG‡)

34. First Step ofHBr Addition to Ethylene • In the addition of HBr the (conceptual) transition-state structure for the first step • The  bond between carbons begins to break • The C–H bond begins to form • The H–Br bond begins to break

35. Reaction Intermediates • If a reaction occurs in more than one step, it must involve species that are neither the reactant nor the final product • These are called reaction intermediates or simply “intermediates” • Each step has its own free energy of activation • The complete diagram for the reaction shows the free energy changes associated with an intermediate

36. Reaction is thermodynamically favourable when the free energy of products is lower than free energy of reactants

37. Energy Diagrams for Single-step Reaction

38. ALKANES

39. Petroleum – natural source of alkanes Natural gas C1 – C4 Asphalt PETROLEUM Straight-run gasoline C5 – C11 b.p. 30-200C Lubricating oil Waxes Catalytic reforming to aromatics: benzene, toluene Kerosene C11 – C14 b.p. 175-300C Gas oil C14 – C25 b.p. 275-400C Catalytic cracking to C3- C5 Catalytic recombination to C7- C10

40. Alkane isomerism

41. Types of Alkyl Groups • Classified by the connection site • a carbon at the end of a chain (primary alkyl group) • a carbon in the middle of a chain (secondary alkyl group) • a carbon with three carbons attached to it (tertiary alkyl group)

42. Types of HydrogenAtoms

43. Alkane intermolecular forces + - + - + - + - + - + - + - + - + - + - + - + - Attractive van der Waals forces caused by temporary dipoles

44. Alkane intermolecular forces Neopentane Pentane (2,2-dimethylpropane) Straigt-chain versus branched alkanes

45. Physical properties of alkanes C1 – C4gases at standard conditions C5 – C19 liquids at standard conditions > C20 solids at standard conditions

46. Physical properties of alkanes

47. Physical properties of alkanes Straigt-chain versus branched alkanes

48. Chemical behaviour of alkanes 1. Combustion of alkanes (oxidation reaction) 2. Chlorination (bromination) of alkanes

49. Chlorination of alkanes – radical chain reaction mechanism 1st step - initiation

50. Chlorination of alkanes – radical chain reaction mechanism 2nd step - propagation