2. Chapter 2 Molecular Orbitals

3. Problem Assignment Skip Section 2.5 1, 23 a, d, e, f, h, i5a 8 - 1315 16, 17 21 a, c, e, f, h 22 a, e, f, g, h23 a, c, e, f, h 25 a, c, e, g, j27, 28, 29 30 a, b, d, e32 d, e

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

5. H. H. HH.. HH HH bonding antibonding   1s ENERGYENERGY

6. node Two atomic orbitals give two molecular orbitals HH HH bonding antibonding   Antibonding orbitals have one node ( a place where the electron density is zero ). The bonding orbital has no node. 1s

7. 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!! Formation of Molecular Orbitals

8.  -Bond in the hydrogen molecule (Bonding Molecular Orbital

9.    Antibonding Orbital in the hydrogen molecule

10.  -Bonding Molecular Orbital and the   -Antibonding Orbital are on top of each other! However, they have different energies!

11. Sect. 2.3 Sigma and pi bonds: Types of Molecular Orbitals There are three important types of molecular orbitals –sigma  –pi  –n Each of these types of orbitals will be discussed further...

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

13.   2p-2p   2p-3d Some Common  Type Bonds Side-to-Side Overlap

14. Pi (  ) Bonds In a multiple bond, the first bond is a sigma (  bond and the second and third bonds are pi (  bonds.      Pi bonds are formed differently than sigma bonds.

15. Types of Bonds:  and  Non-Bonded Pairs : n.. :  n Standard Symbols

16. Sect 2.4: Bond Energy

17. BOND STRENGTHS - C-H SINGLE BONDS C-H C-H 1.10 A 101 C-H 1.08 A 106 CH 2 =CH 2 C-H 1.06 A 121 HC=CH CH 3 CH 3 bond length per mole measured bond energy molecule Kcal = To convert Kcal to an approximate value for Kj, multiply by 4. The book uses Kj.

18. BOND STRENGTHS - MULTIPLE BONDS CC bond bond energy molecule bond length per mole measured Kcal C-C 1.54 A 83 CH 3 CH 3 C=C 1.34 A 146 CH 2 =CH 2 C=C 1.21 A 198 HC=CH = = To convert Kcal to an approximate value for Kj, multiply by 4. The book uses Kj.

19. 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).

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

21. Basic Shapes of Molecules

22. V ALENCE S HELL E LECTRON P AIR R EPULSION VSEPR THEORY 4 pairtetrahedral109 o 28’sp 3 ( pyrimidal, angular ) 3 pairtrigonal planar120 o sp 2 2 pairlinear180 o sp pairsgeometryangles hybridization

23. TETRAHEDRAL GEOMETRY 4 pairs of electrons = tetrahedral 109 o 28’ sp 3

24. Tetrahedral: Methane

25. 3 pairs = trigonal planar (120 o ) 120 o carbocation all repulsions equal 3 equivalent bonds not a stable molecule, but an “intermediate” ion - it reacts quickly

26. Ethylene is a trigonal planar molecule

27. Linear: Acetylene

28. 2 pair linear ( 180 o ) CH 3 MgCH 3 HCN HCCH CH 3 CCCH 3 CH 2 CCH 2 CH 2 CO both 180 o incomplete octetpi bonded all 180 o triple bonds count as only one pair EXPERIMENTAL RESULTS

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

30. Tetrahedral 4 pairs in the valence shell (no double or triple bonds) sp 3 Hybridization

31. sp 3 sp 3 Hybrid Orbital The hybrid orbital has more density in the bonding lobe than a p orbital and forms stronger bonds. ( cross section ) The shape shown is calculated from quantum theory. To avoid confusion the back lobe is omitted from the diagrams, already shown, and the front lobe is elongated to show its direction. omitted Courtesy of Professor George Gerhold

32. C Carbon has 4 valence electrons, 2s 2 2p 2.... Carbon can form single, double or triple bonds sp 3,sp 2 and sp hybrid orbitals. Let’s do sp 3 first.

33. 2s hybridize sp 3 2s promote 2p

34. C tetrahedral geometry 109 o 28’ FORMATION OF TETRAHEDRAL HYBRID ORBITALS 4 hybrid orbitals sp 3

36. Molecular formula of methane: CH 4 Structural formula of methane: C H H H H Remember, 4 bonds mean sp 3 hybrid orbitals. Hydrogen is unhybridized -- 1s orbital

37. C H H H H C H H H H INDICATING BOND GEOMETRY WITH LINES solid lines in plane dotted or dashed line behind plane ( away from you ) wedged line infront of plane ( toward you ) tetrahedral sp 3

38. Ethane, CH 3 -CH 3

39. C C H H HHH H

40. sp 2 hybridization: double bond

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

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

43. C C Note that a double bond consists of a  and a  type bond

44. CC HHHH

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

46. sp hybridization: triple bond

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

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

49. CC HH

50. This model shows lines for  bonds and p-orbitals for the two p bonds.

51. COMPARISON OF THE HYBRIDS

52. COMPARISON OF SP x HYBRID ORBITALS more “p” character more “s” character sp 3 sp 2 sp bigger “tail” more electron density in the bonding lobe “cusp” Orbital plots courtesy of Professor George Gerhold 25% s-character 33% s-character50% s-character

53. WHY DO HYBRIDS FORM? 2. They form stronger bonds. 3. They have better directionality for forming bonds. 1. Electron pair repulsions are minimized. 4. They allow more bonds to form. All of these factors make a lower energy for the molecule.

54. Sections 2.8 - 2.10 Been there done that! Alkanessp 3 Alkenessp 2 Alkynessp

55. Bond Lengths and Strengths revisited Shorter = Stronger = (more) S-character

56. BOND STRENGTHS - C-H SINGLE BONDS C-H C-H sp 3 -1s 1.10 A 101 C-H sp 2 -1s 1.08 A 106 CH 2 =CH 2 C-H sp-1s 1.06 A 121 HC=CH CH 3 CH 3 NOTE : more S = Shorter = Stronger Hybridization makes stronger bonds. bond type length per mole measured bond energy molecule Kcal = increasing s-character

57. BOND STRENGTHS - MULTIPLE BONDS CC bond bond bond energy molecule bond type length per mole measured Kcal C-C sp 3 -sp 3 1.54 A 83 CH 3 CH 3 C=C sp 2 -sp 2 1.34 A 146 CH 2 =CH 2 C=C sp - sp 1.21 A 198 HC=CH = = increasing s-character shorter NOTE : more S = Shorter = Stronger

58. Individual Bond Strengths in Multiple Bonds CC bond bond bond energy individual bond type length  bond strengths Kcal C-C sp 3 -sp 3 1.54 A 83  C-C 83 C=C sp 2 -sp 2 1.34 A 146 C=C sp - sp 1.21 A 198 = Typical  bonds have a bond energy of about 50 kcal/mole p-p 2 p-p  C-C 96  C-C 50  C-C 98  C-C 50 each Kcal

59.  bonds are weaker than sigma (  ) bonds

60. Sect. 2.11 Summary of hydridization of carbon atoms

61. CONSTRUCTION BLOCKS THE HYBRIDS ARE “MOLECULAR LEGOS” EACH IS USED IN A SPECIFIC BONDING SITUATION

62. X Y Z TETRAHEDRAL TRIGONAL PLANAR LINEAR 4 orbitals3 orbitals2 orbitals sp 3 sp 2 sp HYBRID CONSTRUCTION BLOCKS 109 o 28’ 120 o 180 o

63. Shapes of organic molecules and orbital hybridization Single bonds Double bonds Triple bonds “Twisted” molecules

64.  -Bonded Molecules: Single bonds

65. C C H H HHH H Book keeping 6 C-H  bonds sp 3 - 1s 1 C-C  bond sp 3 - sp 3 ethane

66. O C H HH H Methanol Book keeping 3 C-H  bonds sp 3 - 1s 1 C-O  bond sp 3 - sp 3 1 O-H  bond sp 3 - s Oxygen is sp 3 hydridized

67. Now Look at Some  -Bonded Molecules: Double bonds

68. sp 2 p CC H H H H used for  bond used for  bonds Ethylene

69. This model of ethylene shows lines for  bonds and p-orbitals for the pi bond. The second structure shows the pi bonding molecular orbital picture. Book keeping 4 C-H  bonds sp 2 - 1s 1 C-C  bond sp 2 - sp 2 1 C-C  bond p - p

70. sp 2 CO HH used for  bonds Book keeping 2 C-H  bonds sp 2 - 1s 1 C-O  bond sp 2 - sp 2 1 C-O  bond p - p Oxygen is hybridized sp 2 Formaldehyde or Methanal

71. Some  -Bonded Molecules: Triple bonds

72. CC HH Acetylene or Ethyne

73. Book keeping 2 C-H  bonds sp - 1s 1 C-C  bond sp - sp 2 C-C  bond p - p

74. NC H sp

75. This model shows lines for  bonds and p-orbitals for the two p bonds. Book keeping 1 C-H  bonds sp - 1s 1 C-N  bond sp - sp 2 C-N  bond p - p Nitrogen is sp hybridized

76. A  -Bonded Molecule that is “twisted”

77. CCC H H H H allene sp 2 sp C H H HH end view molecule has a twist in the center

78. CCC H H H H sp 2 sp Book keeping 4 C-H  bonds sp 2 - 1s 2 C-C  bond sp 2 - sp 2 C-C  bond p - p Allene

79. Predict Angles, assuming no distortions amide azo compound Remember this functional group? Not a functional group you have to know right now... :

80. Sect. 2.12 p  -d  Backbonding

81. .. : : : : : : - - +2 p  -d  BACKBONDING 3d 2p  -bond S O S O.. :: - : :.. empty 3d orbital “backbonding” Lewis Diagram Expanded Octet Structure no charges MECHANISM OF sulfur has 3d orbitals

82. Section 2.13 Molecular Distortions: VSEPR Revisited Five situations: 1) electron pair repulsion 2) double bond 3) double bond and electronegativity 4) effect of electronegative atoms 5) steric repulsion

83. C H H H H : :.. symmetrical molecule all repulsions are equal perfect tetrahedral all angles 109 o 28” N H H :.. H angle becomes larger repulsion smaller repulsion angle becomes smaller not all pairs are equivalent the unshared pairs repel more strongly than the bonded pairs 1) Electron Pair Repulsion 107 o H-N-H angle reduced to 107 o 109 o 28”

84. H O H H C H H H H N H H 109 o 28’ 107 o 105 o tetrahedral pyrimidal bent Perfect Symmetry 4 pair = tetrahedral ( 109 o 28’) 1) Electron Pair Repulsion

85. C...... larger repulsion one pair : two pairs smaller repulsion one pair : one pair > 120 o < 120 o In alkenes the C=C H angle is typically larger - than the H C H angle. -- 2) Double bond C H H

86. 3) Double bond and Electronegativity more repulsion less repulsion Electrons in C-H bonds are shared nearly equally. Electrons in a C-Cl bond are closer to chlorine and further from each other. POLAR BOND The electronegative Cl draws electrons closer to that end of the bond. NONPOLAR BOND Electrons in C-H bonds are closer to carbon (near center of bond) than in the case below. C...... H H r.. C.. Cl polar R Fluorines cause even smaller angles. CCH 2 H H 121.5 o 117 o CCH 2 Cl 123 o 114 o CCH 2 F F 125 o 110 o

87. 3) Double bond and electronegativity 117 o 121.5 o 123 o 125 o 117 o 114 o 110 o 116 o 122 o 116 o 124.5 o 126 o 111 o 108 o 122 o

88. 4) Effect of Electronegative atoms

89. C CH 3 H H :.. 5) Steric Repulsions.. “Steric Repulsion” 112 o 106 o The CH 3 groups are so large that they push against each other in space, opening the angle.

90. Sect. 2.14: small ring compounds cyclopropane cyclobutane

91. Sect. 2.15 Orbital Array Diagrams 1) Show the types of orbitals present in a molecule: bonding antibonding nonbonding 2) Shows which orbitals are occupied by electrons

92. 0     n ENERGYENERGY -H-H  +H+H  RELATIVE ENERGIES OF THE MOLECULAR ORBITALS IN THE ORBITAL ARRAY OF A MOLECULE IN THE ORBITAL ARRAY OF A MOLECULE note vertical symmetry* about zero bond energy * symmetry is only approximate

93.  bonds C-H 98 C-C 88 CHCH  CC  CC  CH ENERGYENERGY 14 valence e - ORBITAL ARRAY DIAGRAM FOR ETHANE 0 Kcal  all bonds are   -98 -88 KCALKCAL energy released when bond is formed, stronger bonds release more energy MOLEMOLE

94. C H CC CC CC C H C H **  **  C-H bonding and antibonding orbitals C-C bonding and antibonding orbitals sp 3 1s sp 3

95. Most pi (  ) bonds have a bond energy of 50 Kcal / mole MOLECULES WITH PI BONDS When the total energy of a multiple bond is given, you must subtract the energy of the pi bonds to obtain the sigma bond energy. C=C 146 Kcal/mole C-C = 96 Kcal/mole ( 146 - 50 = 96 ) both bonds thus: TOTAL BOND ENERGY

96.  C-H 106 C-C 96* CHCH  CC  CC  CH ENERGYENERGY 12 valence e - ORBITAL ARRAY DIAGRAM FOR ETHYLENE 0 Kcal  CC  CC  C=C 146 *est at (146-50) = 96    -50 -95 -106

97.  C-H 121 C-C 98* CHCH  CC  CC  CH ENERGYENERGY 10 valence e - ORBITAL ARRAY DIAGRAM FOR ACETYLENE 0 Kcal  CC  CC 2  C C 198 *est at (198-100) = 98    = = -50 -98 -121

98. NONBONDING ORBITALS Nonbonding orbitals stay on their original atom during bonding…. They do not form an antibonding orbital …. and their energy changes very little (unshared pairs)

99.  C-H 87 C-O 126* CHCH  CO  CO  CH ENERGYENERGY 12 valence e - ORBITAL ARRAY DIAGRAM FOR FORMALDEHYDE 0 Kcal  CO  CO  C=O 176 *est at (176-50) = 126   nOnO n 

100. Test Topics: Chapter 1 You should work through the sample test! Lewis structures (Section 1.5) Remember: not all atoms have octets! Polar covalent bonds (Section 1.8) Formal charges (Section 1.9) Isomers (Section 1.13) Classes of compounds (Sect. 1.14)

101. Test Topics: Chapter 2 Sigma and pi bonds (Section 2.2 and 2.3) Antibonding orbitals, n-electrons (Sect. 2.3) VESPR and shapes of molecules (Sect. 2.6) sp 3, sp 2 and sp orbitals (Sect. 2.7-2.10) Orbital “bookeeping” and angles associated with various molecules (Sect. 2.11)

102. Chapter 2, continued Orbital “bookeeping,” comparison of bond strengths for each type (Sect. 2.11) VSEPR revisited (distortions) (Sect. 2.13) Small rings (Section 2.14) Molecular orbital arrays (Section 2.15)