Organic Reaction 1

5. An Overview of Organic Reactions

  • Based on McMurry's Organic Chemistry, 7th edition.

2. Why this chapter?

  • Understanding organic and/or biochemistry requires:

    • What occurs

    • Why and how chemical reactions take place

  • The chapter will describe how a reaction can be analyzed.

3. 5.1 Kinds of Organic Reactions

  • Organic reactions typically observed through:

    • What occurs

    • Learning how it happens

  • Common Patterns:

    • Addition Reactions: Two molecules combine.

    • Elimination Reactions: One molecule splits into two.

Substitution Reactions

  • Parts from two molecules exchange.

    • E.g., Methane (alkane) and Chlorine react in the presence of light to form Chloromethane (alkyl halide).

Rearrangement Reactions

  • A molecule undergoes changes in atom connectivity.

    • E.g., 1-Butene can rearrange to form 2-Butene via an acid catalyst.

6. 5.2 How Organic Reactions Occur: Mechanisms

  • Mechanisms describe the steps leading from reactants to products in reactions.

  • Observing the transformation gives insight into the underlying mechanisms.

7. Steps in Mechanisms

  • Steps can involve the formation or breaking of covalent bonds.

    • May occur individually or in combination (concerted steps).

8. Types of Steps in Reaction Mechanisms

  • Bond Formation/Breakage:

    • Symmetrical Bonds: Homolytic.

    • Unsymmetrical Bonds: Heterolytic.

9. Indicating Steps in Mechanisms

  • Curved arrows show bond breaking and forming.

  • Half arrowheads for homolytic steps (radical processes).

  • Complete arrowheads for heterolytic steps (polar processes).

10. 5.3 Radical Reactions

  • Less common than polar reactions.

  • Radicals stabilize electron configurations by reacting and forming bonds.

    • Radicals can lead to substitution or addition reactions.

Steps in Radical Reactions

  1. Initiation: Formation of radicals via homolytic cleavage (e.g., Cl2 under light breaks into Cl atoms).

  2. Propagation: Radicals react, producing further radicals (e.g., Cl reacting with methane to form HCl and CH3).

  3. Termination: Two radicals combine to form a stable product.

11. 5.4 Polar Reactions

  • Polar reactions are influenced by local electron distributions due to electronegativity variations.

  • Partial charges develop on atoms in molecules (e.g., O, F, N, Cl are more electronegative than C).

12. Table 5.1 Polarity Patterns

  • Summarizes polarity in common functional groups:

    • Alcohol, Carbonyl, Alkene, Alkyl halide, Amine, Ether, Thiol, Carboxylic acid, Ketone, etc.

14. Polarizability

  • Polarization involves changes in electron distribution due to surrounding electronic environments.

  • Polar reactions occur between high and low electron density regions.

15. Generalized Polar Reactions

  • Electrophile: Electron-poor species (Lewis acid).

  • Nucleophile: Electron-rich species (Lewis base).

  • Reaction indicated by curved arrow from nucleophile to electrophile.

17. 5.5 An Example of a Polar Reaction

  • Addition of HBr to Ethylene:

    • HBr adds across the double bond where electrons are rich (nucleophile).

18. Mechanism of Addition of HBr to Ethylene

  • HBr reacts with ethylene forming a carbocation intermediate, then bromide adds, creating bromoethane.

19. 5.6 Using Curved Arrows in Polar Reaction Mechanisms

  • Curved arrows depict changes in bonding by tracking electron movement.

  • One arrow corresponds to one reaction mechanism step.

22. 5.7 Describing a Reaction: Equilibria, Rates, and Energy Changes

  • Reactions can reach equilibrium by proceeding forwards or backwards.

  • Equilibrium constant Keq represents the ratio of product to reactant concentrations.

    • Keq > 1 indicates product-favored.

    • Keq < 1 indicates reactant-favored.

24. Free Energy and Equilibrium

  • Gibbs free energy influences product-to-reactant ratios.

    • If Keq > 1, the reaction is exergonic (energy released).

    • If Keq < 1, the reaction is endergonic (energy absorbed).

26. Numeric Relationship of Keq and Free Energy Change

  • Relationship expressed as:

    • [ \Delta G^\circ = -RT \ln K_{eq} ]

    • R = 1.987 cal/(K x mol), T = temperature in K.

27. 5.8 Describing a Reaction: Bond Dissociation Energies

  • Bond dissociation energy (D): Energy required to break a bond to yield radicals.

  • The energy depends on bond types.

    • E.g., C-H bond in methane requires 105 kcal/mol to break.

29. 5.9 Describing a Reaction: Energy Diagrams and Transition States

  • Transition state: Highest energy point in a reaction path.

  • Activation energy (DG‡): Energy required to reach the transition state.

32. 5.10 Describing a Reaction: Intermediates

  • In multi-step reactions, species that are neither reactants nor products are intermediates, each with its own free energy of activation.

33. 5.11 Comparison: Biological vs Laboratory Reactions

  • Laboratory: Typically in organic solvents.

  • Biological: Aqueous medium within cells, facilitated by enzymes that lower activation barriers.