1. 1 Chapter 9 Orbitals and Covalent Bond

2. 2 Molecular Orbitals n The overlap of atomic orbitals from separate atoms makes molecular orbitals n Each molecular orbital has room for two electrons n Two types of MO –Sigma (  ) between atoms –Pi (  ) above and below atoms

3. 3 Sigma bonding orbitals n From s orbitals on separate atoms ++ s orbital +++ Sigma bonding molecular orbital

4. 4 Sigma bonding orbitals n From p orbitals on separate atoms p orbital Sigma bonding molecular orbital  

5. 5 Pi bonding orbitals n p orbitals on separate atoms      Pi bonding molecular orbital

6. 6 Sigma and pi bonds n All single bonds are sigma bonds n A double bond is one sigma and one pi bond n A triple bond is one sigma and two pi bonds.

7. 7 Atomic Orbitals Don’t Work n to explain molecular geometry. n In methane, CH 4, the shape s tetrahedral. n The valence electrons of carbon should be two in s, and two in p. n the p orbitals would have to be at right angles. n The atomic orbitals change when making a molecule

8. 8 Hybridization n We blend the s and p orbitals of the valence electrons and end up with the tetrahedral geometry. n We combine one s orbital and 3 p orbitals. n sp 3 hybridization has tetrahedral geometry.

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12. 12 In terms of energy Energy 2p 2s Hybridizationsp 3

13. 13 How we get to hybridization n We know the geometry from experiment. n We know the orbitals of the atom n hybridizing atomic orbitals can explain the geometry. n So if the geometry requires a tetrahedral shape, it is sp 3 hybridized n This includes bent and trigonal pyramidal molecules because one of the sp 3 lobes holds the lone pair.

14. 14 sp 2 hybridization nC2H4nC2H4nC2H4nC2H4 n Double bond acts as one pair. n trigonal planar n Have to end up with three blended orbitals. n Use one s and two p orbitals to make sp 2 orbitals. n Leaves one p orbital perpendicular.

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18. 18 In terms of energy Energy 2p 2s sp 2 Hybridization 2p

19. 19 Where is the P orbital? n Perpendicular The overlap of orbitals makes a sigma bond (  bond) The overlap of orbitals makes a sigma bond (  bond)

20. 20 Two types of Bonds n Sigma bonds from overlap of orbitals. n Between the atoms. Pi bond (  bond) above and below atoms Pi bond (  bond) above and below atoms n Between adjacent p orbitals. n The two bonds of a double bond.

21. 21 CC H H H H

22. 22 sp 2 hybridization n When three things come off atom. n trigonal planar n 120º One  bond,  lp =3 One  bond,  lp =3

23. 23 What about two n When two things come off. n One s and one p hybridize. n linear

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25. 25 sp hybridization n End up with two lobes 180º apart. n p orbitals are at right angles Makes room for two  bonds and two sigma bonds. Makes room for two  bonds and two sigma bonds. n A triple bond or two double bonds.

26. 26 In terms of energy Energy 2p 2s Hybridization sp 2p

27. 27 CO 2 C can make two  and two  C can make two  and two  O can make one  and one  O can make one  and one  COO

28. 28 N2N2N2N2

29. 29 N2N2N2N2

30. 30 Breaking the octet n PCl 5 n The model predicts that we must use the d orbitals. n dsp 3 hybridization n There is some controversy about how involved the d orbitals are.

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32. 32 dsp 3 n Trigonal bipyrimidal can only  bond. can only  bond. can’t  bond. can’t  bond. n basic shape for five things.

33. 33 PCl 5 Can’t tell the hybridization of Cl Assume sp 3 to minimize repulsion of electron pairs.

34. 34 d 2 sp 3 n gets us to six things around n Octahedral n Only σ bond

35. 35 Molecular Orbital Model n Localized Model we have learned explains much about bonding. n It doesn’t deal well with the ideal of resonance, unpaired electrons, and bond energy. n The MO model is a parallel of the atomic orbital, using quantum mechanics. n Each MO can hold two electrons with opposite spins n Square of wave function tells probability

36. 36 What do you get? n Solve the equations for H 2 n H A H B n get two orbitals n MO 2 = 1s A - 1s B n MO 1 = 1s A + 1s B

37. 37 The Molecular Orbital Model The molecular orbitals are centered on a line through the nuclei The molecular orbitals are centered on a line through the nuclei –MO 1 the greatest probability is between the nuclei –MO 2 it is on either side of the nuclei –this shape is called a sigma molecular orbital

38. 38 The Molecular Orbital Model In the molecule only the molecular orbitals exist, the atomic orbitals are gone In the molecule only the molecular orbitals exist, the atomic orbitals are gone MO 1 is lower in energy than the 1s orbitals they came from. MO 1 is lower in energy than the 1s orbitals they came from. –This favors molecule formation –Called an bonding orbital MO 2 is higher in energy MO 2 is higher in energy –This goes against bonding –antibonding orbital

39. 39 The Molecular Orbital Model Energy MO 2 MO 1 ssss ssss H2H2

40. 40 The Molecular Orbital Model We use labels to indicate shapes, and whether the MO’s are bonding or antibonding. We use labels to indicate shapes, and whether the MO’s are bonding or antibonding. –MO 1 =  1s –MO 2 =  1s * (* indicates antibonding) Can write them the same way as atomic orbitals Can write them the same way as atomic orbitals –H 2 =  1s 2

41. 41 The Molecular Orbital Model Each MO can hold two electrons, but they must have opposite spins Each MO can hold two electrons, but they must have opposite spins Orbitals are conserved. Orbitals are conserved. The number of molecular orbitals must equal the number atomic orbitals that are used to make them. The number of molecular orbitals must equal the number atomic orbitals that are used to make them.

42. 42 H2-H2-H2-H2- Energy  1s  1s * ssss ssss

43. 43 Bond Order n The difference between the number of bonding electrons and the number of antibonding electrons divided by two

44. 44 Only outer orbitals bond n The 1s orbital is much smaller than the 2s orbital n When only the 2s orbitals are involved in bonding Don’t use the  1s or  1s * for Li 2 Don’t use the  1s or  1s * for Li 2 Li 2 = (  2s ) 2 Li 2 = (  2s ) 2 n In order to participate in bonds the orbitals must overlap in space.

45. 45 Bonding in Homonuclear Diatomic Molecules n Need to use Homonuclear so that we know the relative energies. n Li 2 - (  2s ) 2 (  2s *) 1 (  2s ) 2 (  2s *) 1 n Be 2 (  2s ) 2 (  2s *) 2 (  2s ) 2 (  2s *) 2 n What about the p orbitals? How do they form orbitals? n Remember that orbitals must be conserved.

46. 46 B2B2B2B2

47. 47 B2B2B2B2  2p *  2p  2p *  2p

48. 48 Expected Energy Diagram Energy 2s 2p  2s  2p *  2p  2s *  2p *  2p

49. 49 B2B2B2B2 Energy 2s 2p

50. 50 B2B2B2B2 (  2s ) 2 (  2s *) 2 (  2p ) 2 (  2s ) 2 (  2s *) 2 (  2p ) 2 n Bond order = (4-2) / 2 n Should be stable. n This assumes there is no interaction between the s and p orbitals. n Hard to believe since they overlap n proof comes from magnetism.

51. 51 Magnetism n Magnetism has to do with electrons. n Paramagnetism attracted by a magnet. –associated with unpaired electrons. n Diamagnetism attracted by a magnet. –associated with paired electrons. n B 2 is paramagnetic.

52. 52 Magnetism The energies of of the  2p and the  2p are reversed by p and s interacting The energies of of the  2p and the  2p are reversed by p and s interacting The  2s and the  2s * are no longer equally spaced. The  2s and the  2s * are no longer equally spaced. n Here’s what it looks like.

53. 53 Correct energy diagram 2s 2p  2s  2p *  2p  2s *  2p *  2p

54. 54 B2B2B2B2 2s 2p  2s  2p *  2p  2s *  2p *  2p

55. 55 Patterns n As bond order increases, bond energy increases. n As bond order increases, bond length decreases. n Supports basis of MO model. n There is not a direct correlation of bond order to bond energy. n O 2 is known to be paramagnetic. n Movie. Movie

56. 56 Magnetism n Ferromagnetic strongly attracted n Paramagnetic weakly attracted –Liquid Oxygen Liquid OxygenLiquid Oxygen n Diamagnetic weakly repelled –Graphite Graphite –Water Frog Water

57. 57 Examples nC2nC2nC2nC2 nN2nN2nN2nN2 nO2nO2nO2nO2 nF2nF2nF2nF2 nP2nP2nP2nP2

58. 58 Heteronuclear Diatomic Species n Simple type has them in the same energy level, so can use the orbitals we already know. n Slight energy differences. n NO

59. 59 NO 2s 2p

60. 60 You try n NO + n CN - n What if they come from completely different orbitals and energy? n HF n Simplify first by assuming that F only uses one if its 2p orbitals. n F holds onto its electrons, so they have low energy

61. 61 1s 2p  

62. 62 Consequences n Paramagnetic n Since 2p is lower in energy, favored by electrons. n Electrons spend time closer to fluorine. n Compatible with polarity and electronegativity.

63. 63 Names n sp orbitals are called the Localized electron model  and  olecular orbital model  and  olecular orbital model n Localized is good for geometry, doesn’t deal well with resonance. seeing  bonds as localized works well seeing  bonds as localized works well It is the  bonds in the resonance structures that can move. It is the  bonds in the resonance structures that can move.

64. 64  delocalized bonding nC6H6nC6H6nC6H6nC6H6 H H H H H H H H H H H H

65. 65 C2H6C2H6C2H6C2H6

66. 66 NO 3 -