1. CHEMICAL BONDING: ORBITALS Chapter 9

2. LOCALIZED ELECTRON (LE) MODEL A review: views a molecule as a collection of atoms bound together by sharing electrons between their atomic orbitals. The arrangement of valence electrons is represented by the Lewis structure and the molecular geometry can be predicted from the VSEPR model.

3. PROBLEMS WITH THEORY 1. Using the 2p and the 2s orbitals from carbon in methane would result in 2 different types of bonds when they overlap with the 1s from hydrogen. [three 2p/1s bonds and one 2s/2p bond] However, experiments show that methane has FOUR IDENTICAL bonds. 2. Since the 3p orbitals occupy the x, y and z- axes, you would expect those overlaps of atomic orbitals to be at bond angles of 90°. All 4 angles are 109.5°.

4. VALENCE BOND THEORY an extension of the LE model—it’s all about hybridization. Two atoms form a bond when both of the following conditions occur: 1. There is orbital overlap between two atoms. 2. A maximum of two electrons, of opposite spin, can be present in the overlapping orbitals.

5. VALENCE BOND THEORY Because of orbital overlap, the pair of electrons is found within a region influenced by both nuclei. Both electrons are attracted to both atomic nuclei and this leads to bonding. As the extent of overlap increases, the strength of the bond increases. The electronic energy drops as the atoms approach each other but increases when they become too close. This means there is an optimum distance, the observed bond distance, at which the total energy is at a minimum.

6. VALENCE BOND THEORY sigma (σ) bond - overlap of two s orbitals or an s and a p orbital or head-to-head p orbitals. Electron density of a sigma bond is greatest along the axis of the bond. maximum overlap - forms the strongest possible bond; two atoms are arranged to give the greatest possible orbital overlap. Tricky with p orbitals since they are directional. hybrid orbitals - a blending of atomic orbitals to create orbitals of intermediate energy.

7. VALENCE BOND THEORY A carbon atom – no bonding.

8. VALENCE BOND THEORY sp 3 Hybridization Once the carbon atom begins to bond with an atom, the atomic orbitals HYBRIDIZE which changes their shape considerably. There is an energy payoff which explains why they behave this way.

9. VALENCE BOND THEORY The original 2s and three 2p atomic orbitals characteristic of a free carbon atom are combined to form a new set of four sp 3 orbitals.

10. VALENCE BOND THEORY Another way to look at it.

11. PRACTICE ONE Describe the bonding in the ammonia molecule using the LE model. A. Draw the Lewis structure B. Determine arrangement of electron pairs. C. Determine hybridized orbitals needed in the bonding.

12. Ammonia molecule with sp 3 hybridization.

13. HELPFUL HINT It is helpful to think of electron pairs as “sites” of electron density that can be occupied by either a lone pair or a shared pair. If there are 4 “sites” then 4 orbitals need to hybridize so use one s and three p’s to make FOUR equivalent orbitals named sppp or sp 3 orbitals [1 s + 3 p = 4 sp3 orbitals].

14. VALENCE BOND THEORY The tetrahedral set of four sp 3 orbitals of the carbon atom are used to share electron pairs with the four 1s orbitals of the hydrogen atoms to form the four equivalent C-H

15. MULTIPLE BONDING Lowers the number of hybridizing orbitals since unhybridized orbitals are necessary to form the pi bonds Pi (π) bonds - come from the sideways overlap of p atomic orbitals; the region above and below the internuclear axis. NEVER occur without a sigma bond first! may form only if unhybridized p orbitals remain on the bonded atoms occur when sp or sp 2 hybridization is present on central atom NOT sp 3 hybridization. Carbon often does this with N, O, P, and S since they have p orbitals

16. VALENCE BOND THEORY sp 2 Hybridization The set of p’s that are unhybridized are vertical. This is the formation of a sp 2 set of orbitals at angles of 120°. Whenever an atom is surrounded by 3 effective e- pairs, a set of sp 2 hybrid orbitals is required. This molecule would contain a double bond like ethylene. The molecule reserves a set of p’s to form the π bond.

17. VALENCE BOND THEORY The hybridization of the s, px, and py atomic orbitals results in the formation of three sp 2 orbitals centered in the xy plane.

18. VALENCE BOND THEORY The hybridization of the s, p x, and p y atomic orbitals results in the formation of three sp 2 orbitals centered in the xy plane

19. VALENCE BOND THEORY sigma bondpi bond

20. VALENCE BOND THEORY carbon #1 carbon #2 Overlapping unhybridized p orbitals that form the pi bond

21. VALENCE BOND THEORY sp Hybridization This is the formation of a sp set of orbitals at 180°. This molecule would contain a triple bond like ethyne or the double-double arrangement in carbon dioxide. The molecule reserves TWO sets of unhybridized p’s to form the 2 π bonds.

22. VALENCE BOND THEORY When one s orbital and one p orbital are hybridized, a set of two sp orbitals orientated at 180 degrees results.

23. VALENCE BOND THEORY At right is a picture of the two unhybridized p’s on the C atom that are about to make a triple bond. The one labeled at the top is in the plane of this page, the other plain p is in a plane perpendicular to this page.

24. VALENCE BOND THEORY Look at the CO 2 Lewis diagram. The carbon has 2 sites of electron density, each occupied by a double bond, and is therefore sp hybridized while the oxygens have 3 sites (2 lone pairs and a double bond). The oxygen’s have sp 2 hybridization.

25. PRACTICE TWO Describe the bonding in the N 2 molecule.

26. VALENCE BOND THEORY dsp 3 Hybridization Hybridization when there are five pairs of electrons around the central atom. This requires a trigonal bipyramidal arrangement which in turn requires the dsp 3 hybridization.

27. VALENCE BOND THEORY A set of dsp 3 hybrid orbitals on phosphorus atom

28. VALENCE BOND THEORY The orbitals used to form the bonds in PCl 5

29. PRACTICE THREE Describe the bonding in I 3 -, the triiodide ion.

30. VALENCE BOND THEORY d 2 sp 3 Hybridization Hybridization when there are six pairs of electrons around the central atom. This requires an octahedral arrangement which in turn requires the d 2 sp 3 hybridization.

31. VALENCE BOND THEORY An octahedral set of d 2 sp 3 orbitals on sulfur

32. VALENCE BOND THEORY An octahedral set of d 2 sp 3 orbitals on xenon

33. PRACTICE FOUR Why is the xenon atom in XeF 4 hybridized?

35. PRACTICE FIVE Draw the Lewis structure for benzene:

36. BENZENE The sigma bond formations The pi bond formations

37. PRACTICE SIX For each of the following molecules or ions, predict the hybridization of each atom, and describe the molecular structure. a. CO b. BF 4 - c. XeF 2

38. PRACTICE SIX a. CO

39. PRACTICE SIX b. BF 4 -

40. PRACTICE SIX c. XeF 2

41. LE MODEL SUMMARY Three distinct steps: 1. Draw the Lewis structure 2. Determine the arrangements of electron pairs using the VSEPR model. 3. Specify the hybrid orbitals needed to accommodate the electron pairs. Remember that the tendency to minimize energy is more important than the maintenance of the characteristics of atoms as they exist in the free state.