1. Chemical Bonding II: Molecular Geometry and Hybridization of Atomic Orbitals Chapter 10 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2. Molecules are 3-dimensional objects => geometry helps us predict physical and chemical properties melting point boiling point density reactivity Lewis structure shows the number of valence electrons on the central atom of the molecule Electron pairs in the valence shell repel each other => Try to get as far away from each other as possible Valence Shell Electron Pair Repulsion Model (VESPR Model)

3. 10.1

4. Predict geometries based on Lewis Structures Systematically go through all possible geometries. Notation: A = Central atom => atom we predict geometry surrounding B = atoms attached to the central atom E = lone pairs on the central atom (cannot forget lone pairs !!!) start simple = = = > more complex First Case: A - B 2 atoms linear

5. Valence shell electron pair repulsion (VSEPR) model: Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs. AB 2 20 Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry 10.1 linear B B B’sE’s

6. Cl Be 2 atoms bonded to central atom 0 lone pairs on central atom 10.1 Be Cl

7. AB 2 20linear Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 3 30 trigonal planar 10.1 B’sE’s

8. 10.1 B F FF

9. AB 2 20linear Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 3 30 trigonal planar AB 4 40 tetrahedral B’sE’s

10. 10.1 C H H H H

11. AB 2 20linear Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 3 30 trigonal planar AB 4 40 tetrahedral AB 5 50 trigonal bipyramidal trigonal bipyramidal B’sE’s

12. 10.1 P F F F F F

13. AB 2 20linear Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 3 30 trigonal planar AB 4 40 tetrahedral AB 5 50 trigonal bipyramidal trigonal bipyramidal AB 6 60 octahedral B’sE’s Think of axis system x y z

14. 10.1 F F F F F F S

16. bonding-pair vs. bonding pair repulsion lone-pair vs. lone pair repulsion lone-pair vs. bonding pair repulsion >> Lone pairs take up more room than bonding pairs. Lone pairs

17. Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 3 30 trigonal planar AB 2 E21 trigonal planar bent B’sE’s 118.5° for SO 2 120°

18. Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 3 E31 AB 4 40 tetrahedral trigonal pyramidal B’sE’s

19. Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 4 40 tetrahedral AB 3 E31tetrahedral trigonal pyramidal AB 2 E 2 22tetrahedral bent H O H B’sE’s

20. Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 5 50 trigonal bipyramidal trigonal bipyramidal AB 4 E41 trigonal bipyramidal distorted tetrahedron B’sE’s

21. Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 5 50 trigonal bipyramidal trigonal bipyramidal AB 4 E41 trigonal bipyramidal distorted tetrahedron AB 3 E 2 32 trigonal bipyramidal T-shaped Cl F F F B’sE’s

22. Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 5 50 trigonal bipyramidal trigonal bipyramidal AB 4 E41 trigonal bipyramidal distorted tetrahedron AB 3 E 2 32 trigonal bipyramidal T-shaped AB 2 E 3 23 trigonal bipyramidal linear I I I B’sE’s

23. Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 6 60 octahedral AB 5 E51 octahedral square pyramidal Br FF FF F B’sE’s

24. Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 6 60 octahedral AB 5 E51 octahedral square pyramidal AB 4 E 2 42 octahedral square planar Xe FF FF B’sE’s

25. 3 4 5 6

26. Predicting Molecular Geometry 1.Draw Lewis structure for molecule. 2.Count number of lone pairs on the central atom and number of atoms bonded to the central atom. 3.Use VSEPR to predict the geometry of the molecule. What are the molecular geometries of SO 2 and SF 4 ? SO O AB 2 E bent S F F F F AB 4 E distorted tetrahedron

27. Dipole Moments and Polar Molecules 10.2 H F electron rich region electron poor region    = Q x r Q is the charge r is the distance between charges 1 D = 3.36 x 10 -30 C m

28. 10.2

30. Which of the following molecules have a dipole moment? H 2 O, CO 2, SO 2, and CH 4 O H H dipole moment polar molecule S O O CO O no dipole moment nonpolar molecule dipole moment polar molecule C H H HH no dipole moment nonpolar molecule Element EN H 2.1 C 2.5 S 2.5 O 3.5 A non-polar molecule can have polar bonds. A polar molecule must have polar bonds.

31. Does CH 2 Cl 2 have a dipole moment? 10.2

33. HH Bond Dissociation Energy Bond Length H2H2 F2F2 436.4 kJ/mole 150.6 kJ/mole 74 pm 142 pm Valence bond theory – bonds are formed by sharing of e - from overlapping atomic orbitals. Overlap Of 2 1s 2 2p How does Lewis theory explain the bonds in H 2 and F 2 ? Sharing of two electrons between the two atoms. 10.3 FF 1s 1 1s 2 2s 2 2p 5

34. Zero Energy is defined as infinite separation of atoms => no interaction As atoms get closer, the electron of one atom is attracted to the nucleus of the other atom.

35. Change in electron density as two hydrogen atoms approach each other. 10.3

36. Valence Bond Theory and NH 3 N – 1s 2 2s 2 2p 3 3 H – 1s 1 If the bonds form from overlap of 3 2p orbitals on nitrogen with the 1s orbital on each hydrogen atom, what would the molecular geometry of NH 3 be? If use the 3 2p orbitals predict 90 0 Actual H-N-H bond angle is 107.3 0 How do we solve this dilemma??

37. Hybridization – mixing of two or more atomic orbitals to form a new set of hybrid orbitals. 1.Mix at least 2 nonequivalent atomic orbitals (e.g. s and p). Hybrid orbitals have very different shape from original atomic orbitals. 2.Number of hybrid orbitals is equal to number of pure atomic orbitals used in the hybridization process. 3.Covalent bonds are formed by: a.Overlap of hybrid orbitals with atomic orbitals b.Overlap of hybrid orbitals with other hybrid orbitals Can not mix only p’s 1 s + 2 p => 3 hybrid orbitals Conservation in number of orbitals. N H CC or N

38. 1 s + 3 p4 sp 3 Not all orbitals are along an axis. Angle between orbitals is 109.5°

39. 10.4

40. Predict correct bond angle Since non-bonding orbitals are larger than bonding orbitals, expect H - N - H angle to be less than 109.5°.

41. Formation of sp Hybrid Orbitals 10.4 1 s + 2 p 2 sp

42. Formation of sp 2 Hybrid Orbitals 1 s + 2 p3 sp 2

43. # of Lone Pairs + # of Bonded Atoms HybridizationExamples 2 3 4 5 6 sp sp 2 sp 3 sp 3 d sp 3 d 2 BeCl 2 BF 3 CH 4, NH 3, H 2 O PCl 5 SF 6 How do I predict the hybridization of the central atom? Count the number of lone pairs AND the number of atoms bonded to the central atom AB 2 AB 3 or AB 2 E AB 4, AB 3 E or AB 2 E 2 AB 5, AB 4 E, AB 3 E 2, or AB 2 E 3 AB 6, AB 5 E, or AB 4 E 2 Allows for expanded octet

45. NOTE: Not all orbitals need to be used. How do we get multiple bonds? For carbon, if orbitals are sp 3, we can not get enough overlap between 2 carbon atoms. => do not use all p orbitals for hybridization.

47. sp 2 pzpz pzpz

48. Sigma bond (  ) – electron density between the 2 atoms Pi bond (  ) – electron density above and below plane of nuclei of the bonding atoms sp 2 C = C HH HH

49. All electron density is above or below the plane of the molecule in a π bond.

50. For triple bond, leave 2 p orbitals alone.

51. C HH 180°  π

52. Sigma (  ) and Pi Bonds (  ) Single bond 1 sigma bond Double bond 1 sigma bond and 1 pi bond Triple bond 1 sigma bond and 2 pi bonds How many  and  bonds are in the acetic acid (vinegar) molecule CH 3 COOH? C H H CH O OH  bonds = 6 + 1 = 7  bonds = 1 10.5

53. Carbon Hybridization 2s 2 Ground state 2p 2 Promotion of electron 2s 1 2p 3 All single bonds => sp 3 sp 3 One double bond => sp 2 sp 2 p Two double bonds => sp spp C HH C = C HH HH O = C = O

54. Molecular orbital theory – bonds are formed from interaction of atomic orbitals to form molecular orbitals. O O No unpaired e - Should be diamagnetic Experiments show O 2 is paramagnetic 10.6

55. Remember that electrons have a wave nature. ConstructiveDestructive

56. Energy levels of bonding and antibonding molecular orbitals in hydrogen (H 2 ). A bonding molecular orbital has lower energy and greater stability than the atomic orbitals from which it was formed. An antibonding molecular orbital has higher energy and lower stability than the atomic orbitals from which it was formed. 10.6

59. Same configuration of orbitals => fill with allowed number of electrons.

60. 1.The number of molecular orbitals (MOs) formed is always equal to the number of atomic orbitals combined. 2.The more stable the bonding MO, the less stable the corresponding antibonding MO. 3.The filling of MOs proceeds from low to high energies. 4.Each MO can accommodate up to two electrons. 5.Use Hund’s rule when adding electrons to MOs of the same energy. 6.The number of electrons in the MOs is equal to the sum of all the electrons on the bonding atoms. 10.7 Molecular Orbital (MO) Configurations Keep same number of orbitals. Energy shifts have same magnitude. Same as before (Ch 7).

61. bond order = 1 2 Number of electrons in bonding MOs Number of electrons in antibonding MOs ( - ) bond order ½10½ Anti-bonding orbital Bonding orbital No He 2 molecule !!

62. 10.7

63. Delocalized molecular orbitals are not confined between two adjacent bonding atoms, but actually extend over three or more atoms. C C C CC C H H H H H H C C C CC C H H H H H H Benzene – C 6 H 6 Kekulé structures

64. Electron density above and below the plane of the benzene molecule. 10.8