Structure and Bonding of Organic Molecules

1. Structure and Bonding of Organic Molecules Orbital Theory Hybridization and Geometry Polarity Functional Groups

3. Orbitals are Probabilities

4. Bonding in H2TheSigma (s) Bond

5. Wave Depiction of H2

6. Molecular OrbitalsMathematical Combination of Atomic Orbitals

7. Antibonding Molecular OrbitalDestructive Overlap Creates Node

8. Sigma Bonding • Electron density lies between the nuclei. • A bond may be formed by s—s, p—p, s—p, or hybridized orbital overlaps. • The bonding molecular orbital (MO) is lower in energy than the original atomic orbitals. • The antibonding MO is higher in energy than the atomic orbitals. Chapter 2 8

9. s and s* of H2

10. Electron Configurations

11. Ground State Electron Configurations

12. px, py, pz

13. Electron Configuration of Carbon

14. For CARBON, In the Ground State2 bonding sites, 1 lone pair

15. sp3 Hybridization 4 Regions of electron Density link

16. Hybridization of 1 s and 3 p Orbitalsgives 4 sp3 Orbitals

17. sp3 is Tetrahedral GeometryMethane

18. Tetrahderal Geometry

19. Methane Representations

20. Ammonia Tetrahedral GeometryPyramidal Shape

21. All Have the Same GeometryAll Have 4 Regions of Electron DensityAll are sp3 Hybridized

22. Orbital Depiction of Ethane, C2H6 , the s bond

23. Rotation of Single Bonds • Ethane is composed of two methyl groups bonded by the overlap of their sp3 hybrid orbitals. • There is free rotation along single bonds. Chapter 2 23

25. A Model of a Saturated Hydrocarbon

26. sp2 Hybridization3 Regions of Electron Density

27. Hybridization of 1 s and 2 p Orbitals – sp2

28. An sp2 Hybridized Atom

29. The p bondOverlap of 2 parallel p Orbitals

30. Ethylene CH2=CH2

31. Multiple Bonds • A double bond (two pairs of shared electrons) consists of a sigma bond and a pi bond. • A triple bond (three pairs of shared electrons) consists of a sigma bond and two pi bonds. Chapter 2 31

32. Views of Ethylene, C2H4

33. Rotation Around Double Bonds? • Double bonds cannot rotate. • Compounds that differ in how their substituents are arranged around the double bond can be isolated and separated. Chapter 2 33

34. Isomerism • Molecules that have the same molecular formula but differ in the arrangement of their atoms are called isomers. • Constitutional (or structural) isomers differ in their bonding sequence. • Stereoisomers differ only in the arrangement of the atoms in space. Chapter 2 34

35. Constitutional Isomers • Constitutional isomers have the same chemical formula, but the atoms are connected in a different order. • Constitutional isomers have different properties. • The number of isomers increases rapidly as the number of carbon atoms increases. Chapter 2 35

36. Geometric Isomers: Cis and Trans • Stereoisomers are compounds with the atoms bonded in the same order, but their atoms have different orientations in space. • Cis and trans are examples of geometric stereoisomers; they occur when there is a double bond in the compound. • Since there is no free rotation along the carbon–carbon double bond, the groups on these carbons can point to different places in space. Chapter 2 36

37. Formaldehyde

38. sp Hybridization2 Regions of Electron Density

39. The sp Orbital

40. Acetylene, C2H2, 1 s bond2 perpendicular p bonds

41. Molecular Shapes • Bond angles cannot be explained with simple s and porbitals. • Valence-shell electron-pair repulsion theory (VSEPR) is used to explain the molecular shape of molecules. • Hybridized orbitals are lower in energy because electron pairs are farther apart. Chapter 2 41

42. Solved Problem 2 Borane (BH3) is not stable under normal conditions, but it has been detected at low pressure. (a) Draw the Lewis structure for borane. (b) Draw a diagram of the bonding in this molecule, and label the hybridization of each orbital. (c) Predict the H–B–H bond angle. Solution There are only six valence electrons in borane. Boron has a single bond to each of the three hydrogen atoms. The best bonding orbitals are those that provide the greatest electron density in the bonding region while keeping the three pairs of bonding electrons as far apart as possible. Hybridization of an s orbital with two porbitals gives three sp2 hybrid orbitals directed 120° apart. Overlap of these orbitals with the hydrogen 1sorbitals gives a planar, trigonal molecule. (Note that the small back lobes of the hybrid orbitals have been omitted.) Chapter 2 42

43. Summary of Hybridization and Geometry Chapter 2 43

44. Molecular Dipole Moment • The molecular dipole moment is the vector sum of the bond dipole moments. • Depends on bond polarity and bond angles. • Lone pairs of electrons contribute significantly to the dipole moment. Chapter 2 44

45. Lone Pairs and Dipole Moments Chapter 2 45

46. Intermolecular Forces • Strength of attractions between molecules influences the melting point (m. p.), boiling point (b. p.), and solubility of compounds. • Classification of attractive forces: • Dipole–dipoleforces • London dispersions forces • Hydrogen bonding in molecules with —OH or —NH groups Chapter 2 46

47. Dipole–Dipole Interaction Chapter 2 47

48. London Dispersions Chapter 2 48

49. Effect of Branching on Boiling Point • The long-chain isomer (n-pentane) has the greatest surface area and therefore the highest boiling point. • As the amount of chain branching increases, the molecule becomes more spherical and its surface area decreases. • The most highly branched isomer (neopentane) has the smallest surface area and the lowest boiling point. Chapter 2 49

50. Hydrogen Bonds Chapter 2 50