1. Valence Bond Theory

2. How do bonds form? The valence bond model or atomic orbital model was developed by Linus Pauling in order to explain how atoms come together and form molecules. The model theorizes that a covalent bond forms when two orbitals overlap to produce a new combined orbital containing two electrons of opposite spin. This overlapping results in a decrease in the energy of the atoms forming the bond. The shared electron pair is most likely to be found in the space between the two nuclei of the atoms forming the bonds.

3. Example H 2 The newly combined orbital will contain an electron pair with opposite spin just like a filled atomic orbital.

4. Example HF In hydrogen fluoride the 1s orbital of the H will overlap with the half-filled 2p orbital of the F forming a covalent bond.

5. Other Points on the Valence Bond Theory This theory can also be applied to molecules with more than two atoms such as water. Each covalent bond results in a new combined orbital with two oppositely spinning electrons. In order for atoms to bond according to the valence bond model, the orbitals must have an unpaired electron.

6. Covalent Bonding: Orbitals Hybridization The mixing of atomic orbitals to form special orbitals for bonding. The atoms are responding as needed to give the minimum energy for the molecule.

7. sp 3 Hybridization The experimentally known structure of CH 4 molecule can be explained if we assume that the carbon atom adopts a special set of atomic orbitals. These new orbital are obtained by combining the 2s and the three 2p orbitals of the carbon atom to produce four identically shaped orbital that are oriented toward the corners of a tetrahedron and are used to bond to the hydrogen atoms. Whenever a set of equivalent tetrahedral atomic orbitals is required by an atom, this model assumes that the atom adopts a set of sp 3 orbitals; the atom becomes sp 3 hybridized.

8. Figure 9.5. An Energy-Level Diagram Showing the Formation of Four sp 3 Orbitals

9. Figure 9.2. The Valence Orbitals on a Free Carbon Atom: 2s, 2p x, 2p y, and 2p z

10. Figure 9.3. The Formation of sp 3 Hybrid Orbitals

11. Figure 9.6. Tetrahedral Set of Four sp 3 Orbitals

12. Figure 9.7. The Nitrogen Atom in Ammonia is sp 3 Hybridized

13. Figure 9.9. An Orbital Energy-Level Diagram for sp 2 Hybridization

14. Figure 9.8. The Hybridization of the s, p x, and p y Atomic Orbitals

15. A sigma (  ) bond centers along the internuclear axis.  end-to-end overlap of orbitals A pi (  ) bond occupies the space above and below the internuclear axis.  side- to-side overlap of orbitals

16. Figure 9.12. Sigma and Pi Bonding

17. Figure 9.10. An sp 2 Hybridized C Atom

18. Figure 9.11. The  Bonds in Ethylene

19. Figure 9.13. The Orbitals for C 2 H 4

20. Figure 9.16. The Orbital Energy-Level Diagram for the Formation of sp Hybrid Orbitals on Carbon

21. Figure 9.14. When One s Orbital and One p Orbital are Hybridized, a Set of Two sp Orbitals Oriented at 180 Degrees Results

22. Figure 9.17. The Orbitals of an sp Hybridized Carbon Atom

23. Figure 9.18. The Orbital Arrangement for an sp 2 Hybridized Oxygen Atom

24. Figure 9.15. The Hybrid Orbitals in the CO 2 Molecule

25. Figure 9.19. The Orbitals for CO 2

26. Figure 9.20. The Orbitals for N 2

27. Figure 9.21. A Set of dsp 3 Hybrid Orbitals on a Phosphorus Atom

28. Figure 9.23. An Octahedral Set of d 2 sp 3 Orbitals on a Sulfur Atom

29. Figure 9.24. The Relationship of the Number of Effective Pairs, Their Spatial Arrangement, and the Hybrid Orbital Set Required

30. Figure 9.46. A Benzene Ring

31. Figure 9.47. The Sigma System for Benzene

32. Figure 9.48. The Pi System for Benzene

33. The Localized Electron Model Three Steps:  Draw the Lewis structure(s)  Determine the arrangement of electron pairs (VSEPR model).  Specify the necessary hybrid orbitals.

34. Figure 9.45. The Resonance Structures for O 3 and NO 3 -

35. Paramagnetism  unpaired electrons  attracted to induced magnetic field  much stronger than diamagnetism