1. Theories of Covalent Bonding 11.1 Valence Shell Electron Pair Repulsion Theory 11.2 Valence Bond (VB) Theory and Orbital Hybridization 11.3 Molecular Orbital (MO)Theory and Electron Delocalization

3. What are the orbitals that are involved in bonding?

4. We use Valence Bond Theory:

5. Bonds form when orbitals on atoms overlap.

6. A covalent bond forms when the orbitals of two atoms overlap and the overlap region, which is between the nuclei, is occupied by a pair of electrons.

8. Hydrogen fluoride, HF Fluorine, F2

10. As the amount of overlap increases, the energy of the interaction decreases.

11. At some distance the minimum energy is reached.

12. The minimum energy corresponds to the bonding distance (or bond length).

16. There is no unpaired electron available for bonding.

17. We conclude that the atomic orbitals are not adequate to describe orbitals in molecules.

18. We know that the F-Be-F bond angle is 180 (VSEPR theory).

20. We could promote and electron from the 2s orbital on Be to the 2p orbital to get two unpaired electrons for bonding.

21. BUT the geometry is still not explained.

22. We can solve the problem by allowing the 2s and one 2p orbital on Be to mix or form a hybrid orbital.

26. The sp hybrid orbitals in gaseous BeCl2(continued). Figure 11.2 orbital box diagrams with orbital contours

28. The sp2 hybrid orbitals in BF3. Figure 11.3

29. sp2 and sp3 Hybrid Orbitals

30. The sp3 hybrid orbitals in CH4. Figure 11.4

31. Figure 11.5 The sp3 hybrid orbitals in NH3.

32. Figure 11.5 continued The sp3 hybrid orbitals in H2O.

33. Figure 11.6 The sp3d hybrid orbitals in PCl5.

34. The sp3d2hybrid orbitals in SF6. Figure 11.7

35. Key Points Types of Hybrid Orbitals sp sp2 sp3 sp3d sp3d2 Hybrid Orbitals The number of hybrid orbitals obtained equals the number of atomic orbitals mixed. The type of hybrid orbitals obtained varies with the types of atomic orbitals mixed.

37. Step 1 Step 2 Step 3 Figure 11.8 The conceptual steps from molecular formula to the hybrid orbitals used in bonding. Molecular shape and e- group arrangement Molecular formula Lewis structure Hybrid orbitals

38. PROBLEM: Use partial orbital diagrams to describe mixing of the atomic orbitals of the central atom leads to hybrid orbitals in each of the following: PLAN: Use the Lewis structures to ascertain the arrangement of groups and shape of each molecule. Postulate the hybrid orbitals. Use partial orbital box diagrams to indicate the hybrid for the central atoms. SAMPLE PROBLEM 11.1 Postulating Hybrid Orbitals in a Molecule (a) Methanol, CH3OH (b) Sulfur tetrafluoride, SF4 SOLUTION: (a) CH3OH The groups around C are arranged as a tetrahedron. O also has a tetrahedral arrangement with 2 nonbonding e- pairs.

39. hybridized C atom hybridized O atom single C atom single O atom hybridized S atom S atom SAMPLE PROBLEM 11.1 Postulating Hybrid Orbitals in a Molecule continued (b) SF4 has a seesaw shape with 4 bonding and 1 nonbonding e- pairs.

41. Trigonal bipyramidal electron domain geometries require sp3d hybridization.

42. Octahedral electron domain geometries require sp3d2 hybridization.

47. All single bonds are -bonds.

48. -Bonds: electron density lies above and below the plane of the nuclei.

49. A double bond consists of one -bond and one -bond.

50. A triple bond has one -bond and two -bonds.

52. both C are sp3 hybridized s-sp3 overlaps to  bonds sp3-sp3 overlap to form a  bond relatively even distribution of electron density over all  bonds The  bonds in ethane(C2H6). Figure 11.9

54. both C atoms sp2 hybridized;

56. Prentice Hall © 2003 Chapter 9 Multiple Bonds

59. therefore, the C atoms are sp hybridized;

60. the sp hybrid orbitals form the C-C and C-H -bonds;

61. there are two unhybridized p-orbitals;

62. both unhybridized p-orbitals form the two -bonds;

63. one -bond is above and below the plane of the nuclei;

65. Prentice Hall © 2003 Chapter 9 Multiple Bonds

66. Prentice Hall © 2003 Chapter 9 Multiple Bonds

68. In the case of benzene

69. there are 6 C-C  bonds, 6 C-H  bonds,

70. each C atom is sp2 hybridized,

73. localized between C atoms or

74. delocalized over the entire ring (i.e. the  electrons are shared by all 6 C atoms).

75. Experimentally, all C-C bonds are the same length in benzene.

77. Two electrons between atoms on the same axis as the nuclei are  bonds.

78. -Bonds are always localized.

79. If two atoms share more than one pair of electrons, the second and third pair form -bonds.

81. For these molecules, we use Molecular Orbital (MO) Theory.

83. The Central Themes of MO Theory A molecule is viewed on a quantum mechanical level as a collection of nuclei surrounded by delocalized molecular orbitals. Atomic wave functions are summed to obtain molecular wave functions. If wave functions reinforce each other, a bonding MO is formed (region of high electron density exists between the nuclei). If wave functions cancel each other, an antibonding MO is formed (a node of zero electron density occurs between the nuclei).

84. Amplitudes of wave functions subtracted. Figure 11.14 An analogy between light waves and atomic wave functions. Amplitudes of wave functions added

86. each contain a maximum of two electrons;

87. have definite energies;

88. can be visualized with contour diagrams;

90. one has electron density between nuclei (bonding MO);

91. one has little electron density between nuclei (antibonding MO).

92. MOs resulting from s orbitals are  MOs.

95. The total number of electrons in all atoms are placed in the MOs starting from lowest energy (1s) and ending when you run out of electrons.

96. Note that electrons in MOs have opposite spins.

97. H2 has two bonding electrons.

100. paramagnetism (unpaired electrons in molecule): strong attraction between magnetic field and molecule;

101. diamagnetism (no unpaired electrons in molecule): weak repulsion between magnetic field and molecule.

102. Magnetic behavior is detected by determining the mass of a sample in the presence and absence of magnetic field: