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. Molecular Geometry ___ arrangement of atoms affects _________&_________ properties: MP, BP, density, types of reactions, etc. The geometry must be determined __________________, but we can use the valence-shell electron-pair repulsion model (VSEPR) to predict the geometry of a molecule with much success. Remember that the valence shell is the outermost shell that has electrons. Basic assumption: electron pairs in the valence shell ______________________ each other and thus will remain _______________________ each other as possible.

3. Two general rules: 1. Double bonds can be treated as single bonds (with regard to __________________________). 2. VSEPR can be applied to any resonance structure (if the molecule has resonance structures). The valence-shell electron-pair repulsion model accounts for the _________________________ of electron pairs around a central atom in terms of the ___________________between pairs of electrons. So let’s look at some examples that can serve as patterns for you to remember.

4. 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 _________________ B B

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

6. 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 ________ 10.1

8. 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 AB 4 40 _________________

9. 10.1

10. 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 AB 4 40 tetrahedral AB 5 50 _________

11. 10.1

12. 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 AB 4 40 tetrahedral AB 5 50 trigonal bipyramidal trigonal bipyramidal AB 6 60 _________

13. 10.1

14. ______________ << ____________________________

15. Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 3 30 ________ AB 2 E21________ ________ 10.1

16. 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 _________________ _________ ________ 10.1

17. Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 4 40 tetrahedral 10.1 AB 3 E31tetrahedral trigonal pyramidal AB 2 E 2 22_________ H O H

18. Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR 10.1 AB 5 50 _________ AB 4 E41 _________ _________

19. Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR 10.1 AB 5 50 trigonal bipyramidal trigonal bipyramidal AB 4 E41 trigonal bipyramidal distorted tetrahedron AB 3 E 2 32 _________ Cl F F F

20. Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR 10.1 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 _________ I I I

21. Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR 10.1 AB 6 60 _________ AB 5 E51 _________ ________ _________ Br FF FF F

22. Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR 10.1 AB 6 60 octahedral AB 5 E51 octahedral square pyramidal AB 4 E 2 42 _________ _________ Xe FF FF

23. 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 ___________ S F F F F AB 4 E _________ ____________ 10.1

24. Predicting Molecular Geometry The VSEPR model has some theoretical problems—but chemists use it because it is simple and generally leads to useful (reliable) predictions of molecular geometry. 10.1

25. Dipole Moments and Polar Molecules 10.2 H F electron rich region electron poor region    = Q x r Q is the ________________ r is the ___________________________ 1 D = 3.36 x 10 -30 C m Why are the units of D? What does the “D” stand for?

26. 10.2 ___________ molecules are affected by an electric field (p.302)

27. 10.2

28. Which of the following molecules have a dipole moment? H 2 O, CO 2, SO 2, and CH 4 O H H S O O CO O C H H HH

29. 10.2 Dipole Moments of Some Polar Molecules

30. Dipoles (polar molecules) and Microwaves 10.2

31. Read p.304 Lewis theory does not explain why bonds exist or why bonds are of different lengths and different energies. The Valence Bond (VB) Theory and the Molecular Orbital (MO) Theory improve our understanding.

32. Bond Dissociation EnergyBond Length H2H2 F2F2 436.4 kJ/mole 150.6 kJ/mole 74 pm 142 pm __________________ – 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

33. 10.4

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

35. ____________________________ posits that reacting atoms form a stable molecule when/because the potential energy of the system decreases to a minimum (p.306). How does this apply to polyatomic molecules? (p.307) ___________________________________ – atomic orbitals obtained when two or more nonequivalent orbitals of the same atom combine to form a covalent bond _______________________________ – the mixing of atomic orbitals in an atom (central) to produce a set of hybrid orbitals We’ll repeat this in a minute. Let’s look at an example….

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 three 2p orbitals on nitrogen with the 1s orbital on each hydrogen atom, what would the molecular geometry of NH 3 be? Overlap of three 2p orbitals predict _____ 0 Actual H-N-H bond angle is __________0 ! 10.4

37. ____________ – mixing of two or more atomic orbitals to form a new set of hybrid orbitals. 1.Mix at least __________________ 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 _____________ used in the hybridization process. 3.Covalent bonds are formed by: a.Overlap of ______ orbitals with _______ orbitals b.Overlap of ______orbitals with _____________ orbitals 10.4

40. Predict correct bond angle

41. Carefully consider and study the summary of hybridization on the bottom of page 310 in the text.

42. Formation of sp Hybrid Orbitals 10.4

43. Formation of sp 2 Hybrid Orbitals 10.4

44. # 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 ___________________ AND the number of ___________________________________ 10.4 Compare top of p.312

45. 10.5 p.316

46. 10.5

47. Sigma bond (  ) – electron density between the 2 atoms Pi bond (  ) – electron density above and below plane of nuclei of the bonding atoms 10.5

49. p.317

50. 10.5

51. Sigma (  ) and Pi (  ) Bonds Single bond 1 _________ bond Double bond 1 ______ bond & 1 ___ bond Triple bond 1 _____ bond & 2 ____ bonds How many  and how many  bonds are in the acetic acid molecule CH 3 COOH? C H H CH O OH  bonds = 6 + 1 = 7  bonds = 1 10.5

52. Molecular orbital theory – bonds are formed from interaction of atomic orbitals to form molecular orbitals that are associated with the entire molecule. O O No unpaired e - Should be diamagnetic Experiments show O 2 is paramagnetic 10.6 p.318

53. This suggests a fundamental deficiency in Valence Bond Theory—which brings us to the Molecular Orbital Theory. p.318 MO describes covalent bonds in terms of molecular orbitals, that result from interaction of the atomic orbitals of the bonding atoms and are associated with the entire molecule. MO theory posits the formation of two molecular orbitals: one bonding molecular orbital and one antibonding molecular orbital.

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

55. p.319

56. 10.6 Two possible interactions between two equivalent p orbitals and the corresponding molecular orbitals

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

58. A ___________ bond is almost always a sigma bond and a pi bond. A __________ bond is almost always a sigma bond and two pi bonds. Now, carefully consider the Rules Governing Molecular Electron Configuration and Stability p.322

59. Molecular Orbital (MO) Configurations 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

60. 10.6 Note: this ignores the  2s and the   2s orbitals (pp.324f)

61.   s <    s <   s <    s <   py =   px <  2px <    py =    pz <    px p.324 bond order = 1 2 Number of electrons in bonding MOs Number of electrons in antibonding MOs ( - ) p.322

62. Bond order can be a _____________. A bond order of _____ or a _________ bond order means the bond has no stability…and does not exist. Bond order can be used only _____________________ for comparative purposes. p.323

63. _________ = 1 2 Number of electrons in bonding MOs Number of electrons in antibonding MOs ( - ) 10.7 bond order ½10½

64. 10.7 Properties of Homonuclear Diatomic Molecules of the Second-Period Elements p.325

65. MO theory describes bonding in terms of the combination and rearrangement of atomic orbitals to form orbitals associated with the whole molecule. Bonding molecular orbitals increase electron density between the nuclei and are lower in energy than the atomic orbitals from which they are formed. Antibonding molecular orbitals have a region of zero electron density between the nuclei, and an energy level higher than the atomic orbitals from which they are formed. Molecules are stable if the number of electrons in bonding molecular orbitals is greater than the number of electrons in antibonding molecular orbitals. p.326

66. ______________________________ are not confined between two adjacent bonding atoms, but actually extend over three or more atoms. 10.8

67. __________________________ above and below the plane of the _________________________________. 10.8