1. 1 Chapter 9 Covalent Bonding n Includes following concepts: –Hybridization & Localized Electron Model, – Molecular Orbital Model, n Sigma and Pi bonds –Bond Order, –Paramagnetism and Diamagnetism How Often Does The Topic Appear On AP Exam? 9 out of 75 MC Questions & Every Year On FR

2. 2 Chapter 9 Orbitals and Covalent Bond Potential Energy Well, Sigma, & Pi

3. 3 9.1 Hybridization n Atomic Orbitals don’t work to explain molecular geometry. n In methane, CH 4, the shape is tetrahedral. (x-ray crystallography) n The valence electrons of carbon should be 2 in s, and 2 in p in atomic orbital. n the p orbitals would have to be at right angles in atomic orbital. n The atomic orbitals change when making a molecule

4. 4 Hybridization Intro. to Hybridization Defined as the mixing of native Atomic Orbitals to form special atomic orbitals for bonding.

5. 5 Hybridization n In methane we blend the s and p orbitals of the C valence electrons and end up with the tetrahedral geometry. n We combine one s and 3 p atomic orbitals n sp 3 hybridization has tetrahedral geometry.

6. 6 Hybridization sp 3 Native AOHybrid Orbitals

7. 7 Tetrahedral set of sp 3 bonded to H atoms in CH 4 Fig 9.6 p.393

8. 8 In terms of energy Energy 2p 2s Hybridizationsp 3 Atomic Orbitals Hybrid Orbitals Orbital Energy-Level Diagram

9. 9 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.

10. 10 Hybridization sp 2 n C 2 H 4 Ethylene (Alkenes p.1005) n Double bond acts as one pair. n Shape is trigonal planar n Have to end up with three blended orbitals. n Use 1 s and 2 p AO to make sp 2 orbital. n Leaves one p orbital perpendicular.

11. 11 Hybridization sp 2 n When three effective pairs surround the central atom. n Requires a trigonal planar geometry n 120º bond angle Each C forms 3  bonds and 1  bond Each C forms 3  bonds and 1  bond

12. 12 Hybridization sp 2

13. 13 Hybridization sp 2

14. 14 Where is the other P orbital? n One is perpendicular to plane The overlap of orbitals on the atoms axis make sigma bond (  bond) The overlap of orbitals on the atoms axis make sigma bond (  bond)

15. 15 CC H H H H

16. 16 In terms of energy Atomic Orbitals Energy 2p 2s sp 2 Hybridization 2p Orbital Energy-Level Diagram

17. 17 Molecular Orbitals n The overlap of atomic orbitals from separate atoms makes molecular orbitals n Each molecular orbital (MO) has room for two electrons n Two types of MO –Sigma (  ) between atoms on the nuclear axis –Pi (  ) above and below the atoms nuclear axis

18. 18 Sigma bonding orbitals n From s atomic orbitals on separate atoms ++ s orbital Atom s orbital Atom +++ Sigma bonding molecular orbital +

19. 19 Sigma bonding orbitals n From p atomic orbitals on separate atoms p orbital Atom p orbital Atom Sigma bonding molecular orbital  

20. 20 Pi bonding orbitals n p atomic orbitals on separate atoms      Pi bonding molecular orbital Pi Bonds

21. 21 Sigma and pi bonds n All single bonds are sigma bonds n A double bond is » 1 sigma and 1 pi bond n A triple bond is » 1 sigma and 2 pi bonds.

22. 22 Hybridization sp n When two effective pairs surround the central atom. n One s and one p hybridize. n Shape is linear

23. 23 Hybridization sp n End up with two lobes 180º apart. n p orbitals are at right angles

24. 24 In terms of energy Atomic Orbitals Energy 2p 2s Hybridization sp 2p Orbital Energy-Level Diagram

25. 25 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  bonds

26. 26 Breaking The Octet n PCl 5 n Lewis structure shows that P atom is surrounded by 5 electron pairs n 5 pairs requires trigonal bipyramidal n dsp 3 hybridization forms trigonal bipyramidal arrangement

27. 27 dsp 3 n Trigonal bipyrimidal can only  bond. can only  bond. can’t  bond. can’t  bond. n basic shape for five things.

28. 28 PCl 5 Cl hybridizes sp 3 Tetrahedral with 3 lone pairs and 1  bond

29. 29 d 2 sp 3 n SF 6 n gets us to six things around n Octahedral n Only σ bond

30. 30

31. 31 Molecular Orbital Model Thinkwell: Molecular Orbitals

32. 32 Localized Electron Model Summary n Involves 3 distinct steps, in order: 1.Draw the Lewis dot structures. 2.Determine the shape using VSEPR theory. 3.Specify the Hybrid orbitals needed to accommodate the electron pairs.

33. 33 Molecular Orbital Model (MO) n Localized Model we have learned explains much about bonding. –It doesn’t deal well with the ideas of resonance, unpaired electrons, and bond energy. n The MO model is similar to Atomic Orbital (AO) 1.Each MO can hold two electrons with opposite spins 2.Square of wave function tells probability of electron location.

34. 34 What do you get? n Solve quantum equations for H 2 you get 2 MO represented as  MO 2 = 1s A - 1s B  MO 1 = 1s A + 1s B  1s A & 1s B = s orbital for Hydrogen Atoms  MO 1 & MO 2 = bonding probability MO 1 MO 2

35. 35 The Molecular Orbital Model The MO are centered on a line through the nuclei The MO 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 –MO 1 & MO 2 are called sigma molecular orbital

36. 36 The Molecular Orbital Model In the molecule only the MO exist, the AO are gone In the molecule only the MO exist, the AO 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 –Called an antibonding orbital

37. 37 Molecular Orbital Energy Level Diagram Energy MO 2 MO 1 ssss ssss H2H2 Energy Diagram

38. 38 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) called “sigma 1 star” Can write them the same way as atomic orbitals Can write them the same way as atomic orbitals  H 2 =  1s 2

39. 39 The Molecular Orbital Model Each MO can hold two electrons, but they must have opposite spins, just like the Pauli exclusion principle. Each MO can hold two electrons, but they must have opposite spins, just like the Pauli exclusion principle. Orbitals are conserved. Orbitals are conserved. The number of MO must equal the number AO used to make them. The number of MO must equal the number AO used to make them.

40. 40 Energy  1s  1s * ssss ssss Energy Diagram for Hydride Ion H 2 - MO electrons Original AO electrons

41. 41 Bond Order n The difference between the number of bonding electrons and the number of antibonding electrons divided by two

42. 42 Thinkwell: Application of MO Theory Thinkwell: Application of MO Theory

43. 43 Bonding in Homonuclear Diatomic Molecules n In order to participate in bonds the orbitals must overlap in space The 1s orbital is much smaller than the 2s orbital, 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 Relative size of Li 2 Fig 9.31 1s orbital is much smaller than 2s orbital and will not overlap in space.

44. 44 Bonding in Homonuclear Diatomic Molecules n Homonuclear diatomic molecules composed of identical atoms n Li 2 - (  2s ) 2 (  2s *) 1 (  2s ) 2 (  2s *) 1 n Be 2 (  2s ) 2 (  2s *) 2 (  2s ) 2 (  2s *) 2 n Remember that orbitals must be conserved.

45. 45 B 2 Three mutually perpendicular 2p orbitals Above and to left are 2 p orbitals that overlap in parallel fashion. A pair of p orbitals that overlap head-on

46. 46 B2B2B2B2  2p *  2p  2p *  2p

47. 47 Energy 2s 2p  2s  2p *  2p  2s *  2p *  2p Orbital Energy Level Diagram

48. 48 B2B2B2B2 Energy 2s 2p

49. 49 B2B2B2B2 (  2s ) 2 (  2s *) 2 (  2p ) 2 (  2s ) 2 (  2s *) 2 (  2p ) 2 n Bond order = (4-2) / 2 = 1 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.

50. 50 Magnetism n Magnetism has to do with electrons. n Paramagnetism - substance attracted by a magnet. –associated with unpaired electrons. n Diamagnetism – substance repelled by a magnet. –associated with paired electrons. n B 2 is paramagnetic.

51. 51 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.

52. 52 Correct energy diagram 2s 2p  2s  2p *  2p  2s *  2p *  2p

53. 53 B2B2B2B2 2s 2p  2s  2p *  2p  2s *  2p *  2p

54. 54 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.

55. 55 Magnetism n Ferromagnetic strongly attracted n Paramagnetic weakly attracted –p.411 fig. 9.40 n Diamagnetic weakly repelled n Examples fig 9.39 p.410

56. 56 Heteronuclear Diatomic Species n Using atoms adjacent to each other in same period n Because they are in the same energy level, we can use the orbitals we already know. n Example: NO n Paramagnetic predicted by Energy Diagram

57. 57 NO 2s 2p

58. 58 Names n sp orbitals are called the Localized Elect. Model  and  model  and  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.