Chapter 9

1. Chapter 9 Orbitals and Covalent Bond

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

3. Sigma bonding orbitals • From s orbitals on separate atoms + + + + + + Sigma bondingmolecular orbital s orbital s orbital

4.     Sigma bonding orbitals • From p orbitals on separate atoms p orbital p orbital   Sigma bondingmolecular orbital

5.     Pi bonding orbitals • p orbitals on separate atoms     Pi bondingmolecular orbital

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

7. Atomic Orbitals Don’t Work • to explain molecular geometry. • In methane, CH4, the shape is tetrahedral. • The valence electrons of carbon should be two in s, and two in p. • the p orbitals would have to be at right angles. • The atomic orbitals change when making a molecule

8. Hybridization • We blend the s and p orbitals of the valence electrons and end up with the tetrahedral geometry. • We combine one s orbital and 3 p orbitals. • sp3hybridization has tetrahedral geometry.

11. sp3 In terms of energy 2p Hybridization Energy 2s

12. How we get to hybridization • We know the geometry from experiment. • We know the orbitals of the atom • hybridizing atomic orbitals can explain the geometry. • So if the geometry requires a tetrahedral shape, it is sp3 hybridized • This includes bent and trigonal pyramidal molecules because one of the sp3 lobes holds the lone pair.

13. sp2 hybridization • C2H4 • Double bond acts as one pair. • trigonal planar • Have to end up with three blended orbitals. • Use one s and two p orbitals to make sp2 orbitals. • Leaves one p orbital perpendicular.

16. 2p sp2 Hybridization In terms of energy 2p Energy 2s

17. Where is the P orbital? • Perpendicular • The overlap of orbitals makes a sigma bond (s bond)

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

19. H H C C H H

20. sp2 hybridization • When three things come off atom. • trigonal planar • 120º • One p bond, s + lp =3

21. What about two • When two things come off. • One s and one p hybridize. • linear

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

23. 2p sp Hybridization In terms of energy 2p Energy 2s

24. CO2 • C can make two s and two p • O can make one s and one p O C O

25. N2

26. N2

27. Breaking the octet • PCl5 • The model predicts that we must use the d orbitals. • dsp3 hybridization • There is some controversy about how involved the d orbitals are.

28. dsp3 • Trigonal bipyrimidal • can only s bond. • can’t p bond. • basic shape for five things.

29. PCl5 Can’t tell the hybridization of Cl Assume sp3 to minimize repulsion of electron pairs.

30. d2sp3 • gets us to six things around • Octahedral • Only σ bond

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

32. What do you get? • Solve the equations for H2 • HA HB • get two orbitals • MO2 = 1sA - 1sB • MO1 = 1sA + 1sB

33. The Molecular Orbital Model • The molecular orbitals are centered on a line through the nuclei • MO1 the greatest probability is between the nuclei • MO2 it is on either side of the nuclei • this shape is called a sigma molecular orbital

34. The Molecular Orbital Model • In the molecule only the molecular orbitals exist, the atomic orbitals are gone • MO1 is lower in energy than the 1s orbitals they came from. • This favors molecule formation • Called an bonding orbital • MO2 is higher in energy • This goes against bonding • antibonding orbital

35. The Molecular Orbital Model H2 MO2 Energy 1s 1s MO1

36. The Molecular Orbital Model • We use labels to indicate shapes, and whether the MO’s are bonding or antibonding. • MO1 = s1s • MO2 = s1s* (* indicates antibonding) • Can write them the same way as atomic orbitals • H2 = s1s2

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

38. H2- s1s* Energy 1s 1s s1s

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

40. Only outer orbitals bond • The 1s orbital is much smaller than the 2s orbital • When only the 2s orbitals are involved in bonding • Don’t use the s1s or s1s* for Li2 • Li2 = (s2s)2 • In order to participate in bonds the orbitals must overlap in space.

41. Bonding in Homonuclear Diatomic Molecules • Need to use Homonuclear so that we know the relative energies. • Li2- • (s2s)2 (s2s*)1 • Be2 • (s2s)2 (s2s*)2 • What about the p orbitals? How do they form orbitals? • Remember that orbitals must be conserved.

42. B2

43. B2 s2p* s2p p2p* p2p

44. Expected Energy Diagram s2p* p2p* p2p* 2p 2p p2p p2p s2p Energy s2s* 2s 2s s2s

45. B2 2p 2p Energy 2s 2s

46. B2 • (s2s)2(s2s*)2 (s2p)2 • Bond order = (4-2) / 2 • Should be stable. • This assumes there is no interaction between the s and p orbitals. • Hard to believe since they overlap • proof comes from magnetism.

47. Magnetism • Magnetism has to do with electrons. • Remember that spin is how an electron reacts to a magnetic field • Paramagnetism attracted by a magnet. • associated with unpaired electrons. • Diamagnetism repelled by a magnet. • associated with paired electrons. • B2 is paramagnetic.

48. Magnetism • The energies of of the p2p and the s2p are reversed by p and s interacting • The s2s and the s2s* are no longer equally spaced. • Here’s what it looks like.

49. Correct energy diagram s2p* p2p* p2p* 2p s2p 2p p2p p2p s2s* 2s 2s s2s

50. B2 s2p* p2p* 2p 2p s2p p2p s2s* 2s 2s s2s