Chapter 2

1. Chapter 2 Molecular Orbitals

2. Problem Assignment Skip Section 2.5 1, 2 3 a, d, e, f, h, i 5a 8 - 13 15 16, 17 21 a, c, e, f, h 22 a, e, f, g, h 23 a, c, e, f, h 25 a, c, e, g, j 27, 28, 29 30 a, b, d, e 32 d, e

3. . . H H . . H H H H Covalent Bonding ( orbital overlap model ) separated atoms atoms move closer orbitals overlap bond formation STANDARD PICTURE OF COVALENT BONDING

4. s* H H antibonding . . . . H H H H E N E R G Y bonding s H H 1s 1s

5. s* H H node antibonding Two atomic orbitals give two molecular orbitals Antibonding orbitals have onenode ( a place where the electron density is zero ). The bonding orbital has no node . bonding s H H 1s 1s

6. Formation of Molecular Orbitals You have the same number of molecular orbitals as the number of atomic orbitals which combine to form them. However, some of the new orbitals will be antibonding. In other words, when two hydrogen atomic orbitals combine, you get two molecular orbitals. One is a bonding molecular orbital and the other is antibonding molecular orbital. The electron pair go into the bonding orbital! The antibonding orbital remains empty until excitation occurs!!

7. s-Bond in the hydrogen molecule(Bonding Molecular Orbital

8. s*-Antibonding Orbital in the hydrogen molecule

9. s-Bonding Molecular Orbital and thes*-Antibonding Orbital are on top of each other! However, they have different energies!

10. Sect. 2.3 Sigma and pi bonds: Types of Molecular Orbitals • There are three important types of molecular orbitals • sigma (s) • pi (p) • n • Each of these types of orbitals will be discussed further...

11. not symmetric Sigma and Pi Bonds Sigma Bonds symmetric to rotation about internuclear axis s 1s-2p END-TO-END OVERLAP s 2p-2p Pi Bonds SIDE-TO-SIDE OVERLAP p 2p-2p

12. Some Common p-Type Bonds Side-to-Side Overlap p2p-2p p2p-3d

13. Pi (p) Bonds In a multiple bond, the first bond is a sigma (s) bond and the second and third bonds are pi (p) bonds. p p p s s Pi bonds are formed differently than sigma bonds.

14. Standard Symbols Types of Bonds: s and p .. n : s p Non-Bonded Pairs : n

15. Sect 2.4: Bond Energy

16. BOND STRENGTHS - C-H SINGLE BONDS C-H bond energy molecule bond length per mole measured Kcal C-H 1.10 A 101 CH3CH3 C-H 1.08 A 106 CH2=CH2 = C-H 1.06 A 121 HC=CH To convert Kcal to an approximate value for Kj, multiply by 4. The book uses Kj.

17. BOND STRENGTHS - MULTIPLE BONDS CC bond bond energy molecule bond length per mole measured Kcal C-C 1.54 A 83 CH3CH3 C=C 1.34 A 146 CH2=CH2 = = C=C 1.21 A 198 HC=CH To convert Kcal to an approximate value for Kj, multiply by 4. The book uses Kj.

18. Section 2.5 • The “take-away” lesson here is that pure atomic orbitals don’t work for carbon compounds. • You need to hybridize (“mix” s with p orbitals).

19. Section 2.6 VSEPR Theory Valence-shell electron-pair repulsion theory

20. Basic Shapes of Molecules

21. 4 pair tetrahedral 109o28’ sp3 ( pyrimidal, angular ) 3 pair trigonal planar 120o sp2 2 pair linear 180o sp VALENCESHELL ELECTRON PAIR REPULSION VSEPR THEORY pairsgeometryangleshybridization

22. TETRAHEDRAL GEOMETRY 4 pairs of electrons = tetrahedral 109o28’ sp3

23. Tetrahedral: Methane

24. 3 pairs = trigonal planar (120o) 120o 120o carbocation not a stable molecule, but an “intermediate” ion - it reacts quickly all repulsions equal 3 equivalent bonds

25. Ethylene is a trigonal planar molecule

26. Linear: Acetylene

27. EXPERIMENTAL RESULTS H C C H H C N C H C C C H C H M g C H 3 3 3 3 C H C C H 2 2 C H C O 2 triple bonds count as only one pair 2 pair linear ( 180o) incomplete octet pi bonded both 180o all 180o

28. HOW ARE THE OBSERVED BOND ANGLES ACHIEVED? Sect. 2.7 Hybrid Orbitals

29. sp3 Hybridization Tetrahedral 4 pairs in the valence shell (no double or triple bonds)

30. sp3 Hybrid Orbital ( cross section ) The hybrid orbital has more density in the bonding lobe than a p orbital and forms stronger bonds. sp3 To avoid confusion the back lobe is omitted from the diagrams, already shown, and the front lobe is elongated to show its direction. The shape shown is calculated from quantum theory. Courtesy of Professor George Gerhold omitted

31. . . . . C Carbon has 4 valence electrons, 2s22p2 Carbon can form single, double or triple bonds sp3,sp2 and sp hybrid orbitals. Let’s do sp3 first.

32. sp3 hybridize promote 2s 2s 2p 2p

33. FORMATION OF TETRAHEDRAL HYBRID ORBITALS sp3 109o28’ 4 hybrid orbitals C sp3 sp3 sp3 tetrahedral geometry

35. H H H Structural formula of methane: C H Molecular formula of methane: CH4 Remember, 4 bonds mean sp3 hybrid orbitals. Hydrogen is unhybridized -- 1s orbital

36. INDICATING BOND GEOMETRY WITH LINES dotted or dashed line behind plane ( away from you ) solid lines in plane H H C C H H H H H H tetrahedral sp3 wedged line infront of plane ( toward you )

37. Ethane, CH3-CH3

38. H H H C C H H H

39. sp2 hybridization: double bond

40. H H C C H H 2p 2p hybridize 2s sp2 Multiple Bonds and hybridization Ethylene Each carbon is hybridized sp2 . The hydrogens are 1s. One of the double bonds is sp2 - sp2. The other one is p - p.

41. FORMATION OF TRIGONAL PLANAR HYBRID ORBITALS 2p orbital sp2 3 hybrid orbitals C sp2 sp2 trigonal planar

42. C C Note that a double bond consists of a s and a p type bond

43. H H C C H H

44. This model shows lines for s bonds and p-orbitals for the pi bond. The second structure shows the pi bonding molecular orbital picture.

45. sphybridization: triple bond

46. C C H H 2p 2p hybridize sp 2s What about acetylene? Each carbon atom is sp hybridized. The hydrogens are unhybridized, 1s orbitals. Note that a triple bond consists of a s and 2p bonds. The two p bonds use unhybridized p orbitals.

47. FORMATION OF LINEAR HYBRID ORBITALS 2p orbitals sp 2 hybrid orbitals C sp linear

48. H H C C

49. This model shows lines for s bonds and p-orbitals for the two p bonds.

50. COMPARISON OF THE HYBRIDS