1. Molecular Geometry and Polarization

2. Shapes of Molecules Valence Shell Electron Pair Repulsion Theory (VSEPR) a. Bonded electrons b. Lone Pairs

3. 1. Linear (180 o ) BeH 2 CO 2

4. 2. Trigonal Planar (120 o ) NO 3 -

5. 3. Tetrahedral (109.5 o ) CH 4

8. 4. Trigonal Pyramidal (~107 o ) NH 3

9. 5. Bent (~104.5 o ) H 2 O

10. O HH

12. SO 2

13. 6. Trigonal Bipyramid (120 o, 90 o ) PCl 5

14. 7. Octahedral (90 o ) SF 6

16. Shapes of Molecules Ex: Multiple Bonds: N 2 H 2 CO HCN SO 2

17. Shapes of Molecules

18. Models Activity

22. SO 2 2+ SO 2 2- SO 2 SO 3 SF 3 - PF 4 - XeCl 5 + BrF 4 -

23. Predict the molecular geometry of: SnCl 3 - O 3 SeCl 2 CO 3 2- SF 4 IF 5 ClF 3 ICl 4 -

24. WarmUp ClF 4 - SiCl 3 - SO 2 SCl 4 SeO 3 BrCl 5 BrCl 3

28. Polar Molecules 1. Polar molecule – Overall, the electrons are attracted more to one end of an entire molecule 2. Non-Polar Molecule – The electrons are spread out evenly over the entire molecule  -/  + Partial (not full) charges

29. Examples: H2H2OH2H2O CH 4 H 2 CO

30. H 2 H 2 O CH 4 H 2 CO Electron Density

32. Polar Molecules BeCl 2 NH 3 CO 2 SO 2 SF 6 BCl 3 CH 2 Cl 2

33. SCOCH 3 F BH 2 ClPH 3

34. CHF 3 CH 2 F 2 SO 3 SO 3 2- NF 3 CH 3 CHO

35. Hybrid Orbitals A mixing of the atomic orbitals (s, p, d, f) of the central atom Electrons no longer move in the old orbitals, but in a new pattern

36. BeF 2 Isolated Be1s 2 2s 2 (Note that all Be: electrons are paired) To bond Be must unpair some electrons: Bonded Be1s 2 2s 1 2p 1 Be

37. Be is called an “sp” hybrid. Drawings: Isolated BeBeF 2

40. CH 4 Isolated C 1s 2 2s 2 2p 2 Bonded C 1s 2 2s 1 2p 3

41. Isolated C Bonded C sp 3

42. Effect of Lone Pairs Lone pairs do count towards hybridization Ex: H 2 O

43. Try BF 3

44. Examples CCl 4 NH 3 PF 5 SF 6 XeF 4 BrF 3

45. PH 3 H 2 S SF 5 - SF 4 CO 3 2- HCN BrCl 3 CH 4

46. H 2 S SO 2 SO 2 2- AsCl 5 ClF 3 KrF 4

48. Hybrid Orbitals and Multiple Bonds sigma (  ) bonds – single bonds formed by hybrid orbitals pi (  ) bonds – double or triple bonds, not formed by hybrid orbitalsH H – H C=C:N=N:H One  bondOne  bond plusOne  bond plus one  bondtwo  bonds

49. Consider C 2 H 4 Each C is sp 2 Double bond does not count toward hybridization

51. Consider C 2 H 2 Each C is sp hybridized Two  bonds do not count toward hybridization

52. What is the hybridization and bonding types for H 2 CO? Also, what are the bond angles?

53. What is the hybridization and bonding types for acetonitrile (shown)? Also, what are the bond angles? H H -C -C=N: H

54. Delocalized Bonding Adjacent multiple bonds can overlap. Benzene (C 6 H 6 ) All bond lengths are equal

56. Use hybrid orbital theory to explain why all the bonds in the NO 3 - ion are of equal length

61. 12 a) ~110 o b) BF 3 flat (no lone pair) 21. a) (lin)linb) (tetr)tr. Pyc) (Trig bi)ss d) (oh)ohe) (tetr)tetrf) (lin)lin 22 a) (Tetra) Trig. Pyramid b) (Trig planar), Trig pl c) (Tr. Bipy) Td) (Tetra) Tetra e) (Trig Bipy) linf) (Tetra) Bent 24 a) i) Octa (sq.planar)ii) Tetrahedral iii) Trig Bipyr.(see-saw) b) i) Twoii) Oiii) One c) S or Se d) Xe

62. 26. a) 104.5 o, 120 o b) 109.5 o, 120 o c) 107 o, 104.5 o d) 180 o, 109.5 o 28. 2 LP (NH 2 -, ~109 o ), 1 LP (NH 3, 107 o ), 0 LP (NH 4 +, 109 o ) 30. a) ClO 2 - (~109.5 o, 2LP) NO 2 - (120 o, 1 LP) b) XeF 2 (4 LP around the center) 32. a) Lone Pair on P b) Lone Pair on center O 36. Polar = (b), (c), (e) 38.Ortho and meta 44. Not enough p suborbitals 46. SF 2 = sp 3, SF 4 = sp 3 d

63. 48. a) sp 3 b) spc) sp 2 d) sp 3 d e) sp 3 d 2 52. b) N 2 H 4 (sp 3 ), N 2 (sp)c) N 2 stronger bond 54. a) sp 3 (C-H), sp 2 (C-O)b) 36 vec) 26 ve- d) 2 ve- in doublee) 8 ve- in lone pairs 56. a) 1, 120 o 2, 120 o 3, 105 o b) sp 2, sp 2, sp 3 c) 21  bonds

64. 62. 100. In 2 S(I)[Kr]5s 2 4d 10 InS(II) [Kr]5s 1 4d 10 In 2 S 3 (III) [Kr]4d 10 In(III) is smallest (least mutual electron repul) In(III) has the highest lattice energy 102.a) C 2 H 3 Cl 3 O 2 b) C 2 H 3 Cl 3 O 2 c) Structure CCl 3 CH(OH) 2