The document discusses theories of covalent bonding including valence shell electron pair repulsion theory, valence bond theory, orbital hybridization, and molecular orbital theory. It focuses on how orbitals overlap to form bonds between atoms and the different types of hybrid orbitals that can form based on electron domain geometry. Examples are provided to illustrate sp, sp2, sp3, and other hybrid orbitals as well as sigma and pi bonding including delocalized pi bonding in benzene.

1. Theories of Covalent Bonding11.1 Valence Shell Electron Pair Repulsion Theory11.2 Valence Bond (VB) Theory and Orbital Hybridization11.3 Molecular Orbital (MO)Theory and Electron Delocalization

2. Prentice Hall © 2003Chapter 9Covalent Bonding and Orbital OverlapLewis structures and VSEPR do not explain why a bond forms. How do we account for shape in terms of quantum mechanics?

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.

7. There are two electrons of opposite spin in the orbital overlap.Prentice Hall © 2003Chapter 9Covalent Bonding and Orbital OverlapHydrogen, H2

8. Hydrogen fluoride, HFFluorine, F2

9. Prentice Hall © 2003Chapter 9Covalent Bonding and Orbital OverlapAs two nuclei approach each other their atomic orbitals overlap.

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).

13. As the two atoms get closer, their nuclei begin to repel and the energy increases.Prentice Hall © 2003Chapter 9Covalent Bonding and Orbital OverlapAt the bonding distance, the attractive forces between nuclei and electrons just balance the repulsive forces (nucleus-nucleus, electron-electron).

14. Prentice Hall © 2003Chapter 9Hybrid OrbitalsAtomic orbitals can mix or hybridize in order to adopt an appropriate geometry for bonding.

15. Hybridization is determined by the electron domain geometry.sp Hybrid OrbitalsConsider the BeF2 molecule (experimentally known to exist):Prentice Hall © 2003Chapter 9Hybrid Orbitalssp Hybrid OrbitalsBe has a 1s22s2 electron configuration.

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).

19. We also know that one electron from Be is shared with each one of the unpaired electrons from F.Prentice Hall © 2003Chapter 9Hybrid Orbitalssp Hybrid OrbitalsWe assume that the Be orbitals in the Be-F bond are 180 apart.

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.

23. The hybrid orbital comes from an s and a p orbital and is called an sp hybrid orbital.

24. Prentice Hall © 2003Chapter 9Hybrid Orbitalssp Hybrid OrbitalsThe lobes of sp hybrid orbitals are 180º apart.

25. Since only one of the Be 2p orbitals has been used in hybridization, there are two unhybridized p orbitals remaining on Be.The sp hybrid orbitals in gaseous BeCl2.Figure 11.2atomic orbitalshybrid orbitalsorbital box diagrams

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

28. The sp2 hybrid orbitals in BF3.Figure 11.3

29. sp2 and sp3Hybrid Orbitals

30. The sp3 hybrid orbitals in CH4.Figure 11.4

31. Figure 11.5The sp3 hybrid orbitals in NH3.

32. Figure 11.5 continuedThe sp3 hybrid orbitals in H2O.

33. Figure 11.6The sp3d hybrid orbitals in PCl5.

34. The sp3d2hybrid orbitals in SF6.Figure 11.7

35. Key PointsTypes of Hybrid Orbitalsspsp2sp3sp3dsp3d2Hybrid OrbitalsThe 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 1Step 2Step 3Figure 11.8The conceptual steps from molecular formula to the hybrid orbitals used in bonding.Molecular shape and e- group arrangementMolecular formulaLewis structureHybrid 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 11.1Postulating Hybrid Orbitals in a Molecule(a) Methanol, CH3OH(b) Sulfur tetrafluoride, SF4SOLUTION:(a) CH3OHThe groups around C are arranged as a tetrahedron.O also has a tetrahedral arrangement with 2 nonbonding e- pairs.

39. hybridized C atomhybridized O atomsingle C atomsingle O atomhybridized S atomS atomSAMPLE PROBLEM 11.1Postulating Hybrid Orbitals in a Moleculecontinued(b) SF4 has a seesaw shape with 4 bonding and 1 nonbonding e- pairs.

40. Prentice Hall © 2003Chapter 9Hybrid OrbitalsHybridization Involving d OrbitalsSince there are only three p-orbitals, trigonal bipyramidal and octahedral electron domain geometries must involve d-orbitals.

41. Trigonal bipyramidal electron domain geometries require sp3d hybridization.

42. Octahedral electron domain geometries require sp3d2 hybridization.

43. Note the electron domain geometry from VSEPR theory determines the hybridization.Prentice Hall © 2003Chapter 9Hybrid OrbitalsSummaryDraw the Lewis structure.Determine the electron domain geometry with VSEPR.Specify the hybrid orbitals required for the electron pairs based on the electron domain geometry.

46. Prentice Hall © 2003Chapter 9Multiple Bonds-Bonds: electron density lies on the axis between the nuclei.

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.

51. Often, the p-orbitals involved in -bonding come from unhybridized orbitals.Multiple Bonds

52. both C are sp3 hybridizeds-sp3 overlaps to  bondssp3-sp3 overlap to form a  bondrelatively even distribution of electron density over all  bondsThe  bonds in ethane(C2H6).Figure 11.9

53. Prentice Hall © 2003Chapter 9Multiple BondsEthylene, C2H4, has:one - and one -bond;

54. both C atoms sp2 hybridized;

55. both C atoms with trigonal planar electron pair and molecular geometries.overlap in one position - p overlap - electron densityThe  and  bonds in ethylene (C2H4).Figure 11.10

56. Prentice Hall © 2003Chapter 9Multiple Bonds