1. Organic Chemistry Review Part II

2. Organic Chemistry: Carbon Atom 1. Structural Classifications 2. Atomic Theory 3. Dipoles & Resonance 4. Isomers 5. Functional Groups 6. Organic Reactions

3. Organic Chemistry The chemistry of compounds which contain carbon. Carbon forms more compounds than any other element, except hydrogen.

4. Organic Chemistry Major Concepts 1. Structural Classifications 2. Hybridization 3. Charges of Organic Molecules 4. Dipoles & Dipolar Resonance 5. Isomers 6. Functional Groups 7. Organic Reactions

5. Structural Classification of Carbon Atoms Three main classifications are: 1. Primary Carbons 2. Secondary Carbons 3. Tertiary Carbons 4. Quaternary Carbons

6. Primary Carbons Denoted as 1° carbons. Also called terminal or end carbon atoms. Found at the ends of a straight chains or the branches. Covalently bonded to one carbon atom.

7. Secondary Carbons Denoted as 2° carbons. Covalently bonded to two other carbon atoms.

8. Tertiary Carbons Denoted as 3° carbons. Covalently bonded to three other carbon atoms.

9. Quaternary Carbons Denoted as 4° carbons. Covalently bonded to four other carbon atoms.

10. Definitions Valence Bond Theory: Electrons in a covalent bond reside in a region in which there is overlap of individual atomic orbitals. For example, the covalent bond in molecular methane (CH 4 ) requires the overlap of valence electrons:

11. Definitions Types of valence bond theory overlap:

12. Definitions Valence Shell Electron Pair Repulsion (VSEPR) Electron pairs arrange themselves around an atom in order to minimize repulsions between pairs. Carbon has a valence of four and must have a tetrahedral geometry. In methane, each carbon atom must have a bond angle of 109.5 ⁰. This is the largest bond angle that can be attained between all four bonding pairs at once.

13. Definitions Hybridization: Atomic orbitals modify themselves to meet VESPR geometry and valence bond theory. Three types of hybridization for carbon:

14. Hybridization: Valence Bond Theory

15. Hybridization: VSEPR Geometry

16. Hybridizations

17. In sp 3 hybridization, an electron is promoted from a 2s orbital into a p orbital. The 2s orbital and three 2p orbitals form four hybrid orbitals (sp 3 ). Ground state: 1s 2 2s 2 2P x 1 2P y 1 2P z 0 Excited state: 1s 2 2s 1 2P x 1 2P y 1 2P z 1

18. Hybridizations The overlap of each hybrid orbital with a hydrogen atom results in a sigma bond ( σ bond). Only one σ bond can exist between two atoms.

19. Hybridizations sp 3 hybridization of methane:

20. Hybridizations sp 3 hybridization of ethane:

21. Hybridizations In sp 2 hybridization, the 2s orbital and two of the 2p orbitals form three hybrid orbitals (sp 2 ). The P z orbital of each carbon atom remains unhybridized. These unhybridized P z orbitals overlap with one another to form a π-bond.

22. Hybridizations sp 2 hybridization of ethene:

23. Hybridizations sp 2 hybridization and bond rotation:

24. Hybridizations In sp hybridization, the 2s orbital and one 2p orbital form two hybrid orbitals (sp). The triple bond is actually one σ bond and two π bonds.

25. Hybridizations sp hybridization of ethyne: No free rotation

26. Charges in Organic Molecules

27. Definitions Dipole: The measure of net molecular polarity. Formula: the magnitude of the charge Q times the distance r between the charges. μ = Q × r The larger the difference in electronegativities of the bonded atoms, the larger the dipole moment.

28. Definitions Resonance: Part of the Valence Bond Theory Describes the delocalization of electrons within molecules. Used when Lewis structures for a single molecule cannot describe the actual bond lengths between atoms. Structures are not isomers of the target molecule, since they only differ by the position of delocalized electrons.

29. Definitions Resonance Hybrid: The net sum of valid resonance structures. Several structures represent the overall delocalization of electrons within the molecule. A molecule that has several resonance structures is more stable than one with fewer.

30. Definitions Hyperconjugation: The interaction of the electrons in a sigma bond (usually C–H or C–C) with an adjacent empty (or partially filled) non-bonding p-orbital, antibonding π orbital, or filled π orbital. Only electrons in bonds that are β to the positively charged carbon can stabilize a carbocation by hyperconjugation.

31. Carbon Atom Dipoles Carbon- Halogen Bonds

32. Carbon Atom Dipoles C-O, C-S and C-N Covalent Bonds: δ+δ+ δ+δ+ δ-δ- δ-δ-

33. Dipolar Resonance

37. Hyperconjugation A.K.A "no bond resonance". The delocalization of σ-electrons or lone pair of electrons into adjacent π-orbital or p-orbital. Overlapping of σ-bonding orbital or the orbital containing a lone pair with adjacent π-orbital or p-orbital. An α- carbon next to the π bond, carbocation or free radical should be sp 3 hybridized with at least one hydrogen atom bonded to it.

38. Hyperconjugation Other hydrogens on the methyl group also participate due to free rotation of the C-C bond. There is NO bond between an α-carbon and one of the hydrogen atoms. The hydrogen atom is completely detached from the structure. The C-C bond acquires some double bond character and C=C acquires some single bond character.

39. Hyperconjugation

40. Hyperconjugation: Examples

43. Isomers Compounds that have: The same molecular formula. Similar or different types of structural formulas. Different arrangement of atoms.

44. Isomers: Two main classes are: 1. Structural or constitutional 2. Stereoisomers

45. Structural Isomers Also known as constitutional isomers

47. Stereoisomers a. Configurational Geometric or Diastereomers Optical or Enantiomers b. Conformational or Rotamers

48. Diastereomers

49. Geometric Isomers: Examples

51. Optical Isomers

52. Definitions Chiral Molecules - when a molecule and its mirror image cannot completely overlap. They are non-superimposable mirror images of one another. Dextrorotatory (R, +) - a compound whose solution rotates the plane of polarized light to the right (when looking toward the source of light).

53. Definitions Levorotatory (S, -) - a compound whose solution rotates the plane of polarized light to the left (when looking toward the source of light). Racemic Mixture - a mixture of equal amounts of optical isomers. Because the two isomers rotate the plane of polarized light by the same angle in opposite directions, they cancel each other out and have no net effect.

54. Determining L (S, -) or D (R, +) configuration 1. Rank the four substituents according to the atomic numbers of the atoms bonded directly to the double bonded carbons, from highest (1) to lowest (4).

55. Determining L (S, -) or D (R, +) configuration 2. If two substituents have the same ranking: Look at the next atoms in their substituent chains. List the atoms that are two bonds away from the chiral center according to their atomic number, from highest to lowest. Assign the lower number to the substituent that has the atom with the higher atomic number.

56. Determining L (S) or D (R) configuration If it is still the same atom for both substituents, continue down the list until a difference is found and assign a ranking in the same manner. 3. If a substituent has a double or triple bonds in its chain, it is counted as two or three bonds to the same atom.

57. Determining L (S, -) or D (R, +) configuration 4. Determine whether the ranking defines a clockwise or counterclockwise direction.  If clockwise, the projection is an R configuration.  If counterclockwise, it is an S configuration.

58. Determining L (S, -) or D (R, +) configuration

59. L (S, -) Configuration A common optical isomer for amino acids in Biochemistry

60. Optical Isomers: Examples

61. Summary of Isomers

64. Conformational Isomers Also known as Rotamers Stereoisomers that can be interconverted by the rotation of atoms about a σ-bond.

65. Conformational Isomers

66. Rotamers: Examples