CHAPTER 10

1. CHAPTER 10 ORBITAL HYBRIDIZATION and MOLECULAR ORBITALS Problems 1-15 + all bold numbered problems

3. ORBITALS AND BONDING THEORIES • There are two major theories of bonding • The Valence Bond (VB) Theory • Simple • Explains geometry • An extension of the Atomic Theory • Fails to account for some paramagnetic properties • Fails to explain delocalized bonds • The Molecular Orbital (MO) Theory • More complex • Explains magnetic properties • Explains delocalized bonds

4. Two Theories of Bonding • VALENCE BOND THEORY — Linus Pauling • valence electrons are localized between atoms (or are lone pairs). • half-filled atomic orbitals overlap to form bonds

5. Maximum Attraction Minimum Attraction Free atom Repulsion

6. H H ­ ¯ + ­ ­ • • • • sigma bond ( s ) Sigma Bond Formation by Atomic Orbital Overlap Two s orbitals overlap

7. Sigma Bond Formation by Atomic Orbital Overlap Two s orbitals overlap Two p atomic orbitals overlap s & p atomic orbitals overlap H F H + F

8. VALENCE BOND THEORY • Bonds form when: atomic orbitals overlap, and two electrons with opposite spin are present. • The resulting lower energy state is called a covalent bond. • If the overlapping orbitals are 1s type orbitals, the resulting bondis called a s1s + 1s (previous H2) • If the overlapping orbitals are 1s and 2p type orbitals, the resulting bondis called a s1s + 2p. • If the overlapping orbitals are 2p type orbitals, the resulting bondis called a s2p + 2p

9. VALENCE BOND THEORY • VBT works well for explaining atomic orbital overlap between 1s-1s, 2p-2p and 1s-2p to make new bonds • What does VBT say about more complex molecules with more then 2 atom systems

10. VALENCE BOND THEORY • VBT can not account for a 109.5º angle in CH4 • due to 2p orbitals constrained 90º bond angle • AO of 90º do not correspond to 109.5º ?

11. A new theory is needed Linus Pauling proposes Hybrid Orbital Theory

12. Advanced Theories of Chemical Bonding Atomic Orbitals Molecules

13. Hybrid Orbitals • Atomic orbitals in the free atoms are “Mixed” to make a new Hybrid Orbital for molecules with more then 2 atoms • An example of mixing an s and p atomic orbital to make a NEW Hybrid sp orbital • # AO = # of HO

14. New HO from AO + s p3 s p3 s p p s p 3d s p 3d p d sp2 sp3 sp3d + s p2 s p2 s p2 s p p s p3 s p3 p + s p 3d s p 3d s p p s p 3d

15. s p 3d s p 3d2 p d + sp3d2 s p 3d2 s p 3d2 s p p s p 3d2 s p 3d2 d • These new HO do two things • Allow the orbitals to physically rearrange in such a fashion to adopt the 109.5º bond angle in the sp3 for example • Account for 6 coordinate molecules as in the case of sp3d2

16. Hybrid Orbitals: a “Mixing” of AO’s • New Name • sp • sp2 • sp3 • sp3d • sp3d2 • Other possibilities • one s with one p (just described) • one s with two p’s • one s with three p’s • one s with three p’s and one d • one s with three p’s and two d’s These are the newly mixed Hybrid Orbitals

17. Hybridization of Atomic Orbitals Sp3 • To form polyatomic molecules having three or more atoms, it is frequently necessary (always in the second period) to hybridize the atomic orbitals of the central atom to permit maximum orbital overlap and to maximize the number of bonds formed. • Carbon, for example, always forms four (4) bonds.

18. Orbitals needed to hybridize

20. Bonding in CH4 Need to use 4 atomic orbitals — s, px, py, and pz — to form 4 new hybrid orbitals pointing in the correct direction.

21. Bonding in a Tetrahedron Formation of Hybrid Atomic Orbitals 4 C atom orbitals hybridize to form four equivalent sp3 hybrid atomic orbitals.

22. sp3 • These hybrid atomic orbitals are linear combinations of the atomic orbitals. • The number of hybrid orbitals produced is always equal to the number of atomic orbitals hybridized.

23. New Linear Combination

24. Bonding in a TetrahedronFormation of Hybrid Atomic Orbitals 4 C atom orbitals hybridize to form four equivalent sp3 hybrid atomic orbitals.

25. sp3 • Hybridization for water and ammonia is explained on page 445. • Figures 10.8 & 10.9. • (Examples 10.1 and 10.2 page 446.) • Predict the hybridization and geometry for CHCl3 and OF2. Describe the sigma bonds in each molecule.

26. Figure 10.8

27. Figure 10.9

28. sp2 • If one s orbital and two p orbitals are combined, the resulting hybrid orbitals are called sp2 orbitals. • The trigonal planar shape of this set of three orbitals is consistent with the VSEPR theory. • Figure 10.10, page 448, and Figure 10.8 (3rd ed.) illustrates this concept. • The orbital box diagrams can be used to confirm the bonding. • BH3 is and example (show model).

29. Figure 10.10

30. •• • • F • • Boron configuration B ­ ­¯ ­¯ •• •• • • F F • 2p 1s 2s • •• •• Using VB Theory Bonding in BF3 planar triangle angle = 120o

31. Bonding in BF3 • How to account for 3 bonds 120o apart using a spherical s orbital and p orbitals that are 90o apart? • Pauling said to modify VP approach with ORBITAL HYBRIDIZATION • — mix available orbitals to form a new set of orbitals — HYBRID ORBITALS — that will give the maximum overlap in the correct geometry.

32. ­ ­¯ 2p 2s hydridize orbs. rearrange electrons ­ ­ ­ unused p 2 three sp orbital hybrid orbitals Bonding in BF3

33. ­ three ­ ­¯ 120o 2 p 2 s 2 sp hybrid hydridize orbs. rearrange electrons • orbitals ­ ­ ­ ­ ­ unused p 2 three sp orbital hybrid orbitals Bonding in BF3 • The three hybrid orbitals are made from 1 s orbital and 2 p orbitals  3 sp2 hybrids. • Now we have 3, half-filled HYBRID orbitals that can be used to form B- F sigma bonds.

34. ­ three ­ ­¯ 120o 2 p 2 s 2 sp hybrid hydridize orbs. rearrange electrons • orbitals ­ ­ ­ ­ ­ unused p 2 three sp orbital hybrid orbitals Bonding in BF3 An orbital from each F overlaps one of the sp2 hybrids to form a B-F  bond.

35. sp • Figure 10.11, page 448, & Figure 10.9 (3rd ed.) illustrates the outcome of hybridizing an s and a p orbital to produce two sp hybrid atomic orbitals with a linear geometry. • BeF2 is an example of this type hybridization. • Predict the hybridization and geometry for PF3 and BeH2. Describe the sigma bonds in each molecule. • We would expect the first to be sp3 hybridized, but it actually uses the 3p atomic orbital (3rd period) instead and has approximately 90o bond angles.

36. Figure 10.11 36

37. Orbital Hybridization BONDS SHAPE HYBRID REMAIN 2 linear sp 2 p’s 3 trigonal sp2 1 p planar 4 tetrahedral sp3 none

38. dsp3 and d2sp3 • These hybridizations require five (5) and six (6) atomic orbitals respectively and produce an equal number of hybrid atomic orbitals. The electronic geometries are trigonal-bipyramidal, and octahedral. • Figure 10.12provides an overview of these cases and a review of the others. (page 450) • Examples 10.3 & 4 cover PF5 and others. • Describe the hybridization, bonding, and geometry for SF4 and SF6.

39. Figure 10.12

40. Multiple Bonding • The second type of covalent bond is the p bond. • The pi bond (p) results from the sideways overlap of two atomic p orbitals. The electron density is above and below the plane. • The second bond between two atoms is always a p bond as is the third bond in the case of a triple bond. • Figures 11.8 (2nd ed.) & 10.13, 10.14, 10.16 and models.

41. Figure 10.13

42. Figure 10.14

43. Figure 10.16 2 p bonds

44. Bonding in Glycine

45. H H 2 sp 120o C C H H Multiple Bonds Consider ethylene, C2H4

46. H H 2 sp 120o C C H H Sigma Bonds in C2H4

47. ­ ­ ­ ­ ­¯ ­ ­ p 2 2s 2p 3 sp orb. hybrid for p orbitals bond p Bonding in C2H4 The unused p orbital on each C atom contains an electron and this p orbital overlaps the p orbital on the neighboring atom to form the p bond. (See Fig. 10.13)

48. Multiple Bondingin C2H4

49. Consequences of Multiple Bonding Restricted rotation around C=C bond.

50. Multiple Bonding • Describe the hybridization, bonding, and geometry for H2CO and HCN. Also do a valence bond bonding diagram for each. • Figure 10.15 for H2CO bonding. • O3, ozone, is an excellent example of pi resonance and a need for a better theory, the delocalized molecular theory.