1. BA + -- (Bonding) 2p z Atomic Orbitals of Atoms A,B zz AB 2p - - ++ Molecular Orbitals A - - ++ B  *(Antibonding) 2p z Nodal Plane 2p z -2p z -- ++ AB 2p z (A) + 2p z (B) 2p z (A) - 2p z (B) z Axis is taken to be line through A, B z axis z axis Bonding with 2p Orbitals

2. Atomic Orbitals of Atoms A,B Molecular Orbitals AB + + - - 2p x -2p x AB ++ -- 2p x x  (bonding) AB + - AB + + - - (antibonding) ** 2p Nodal Plane 2p x (A) - 2p x (B) 2p x (A) + 2p x (B) z x x z Note: orbital is  to bond (z) axis A,B Nodal Plane Bonding with 2p Orbitals (cont)

3. Notational Detail Oxtoby uses two different notations for orbitals in the 4th and 5th editions of the class text: 2px* in the 4th edition becomes g2px * in the 5th edition 2px in the 4th edition becomes u2px in the 5th edition 2pz in the 4th edition becomes   g2pz in the 5th edition 2pz* in the 4th edition becomes   u2pz * in the 5th edition

4. CORRELATION DIAGRAM f or second period diatomic molecules E ** 2p z ** y ** x  x  y  z ** z ** x ** y  z  x  y 2s **  2p 2s **  (6 Valence electrons for each atom) (3 Valence electrons for each atom) B 2 Z ≤ 7 O 2 Z ≥ 8 Molecular OrbitalsMolecular Orbitals Note unpaired electrons Note order of filling  2p z Degenerate  orbitals

5. Molecular Orbitals of Homonuclear Diatomic Molecules (Å) He 2 Li 2 Be 2 B2B2B2B2 C2C2C2C2 N2N2N2N2 O2O2O2O2 F2F2F2F2 Ne 2 2 4 2 4 6 8 10 12 14 16 1 0 1 0 1 2 3 2 1 0 0.74 52 2.67 2.45 1.59 1.24 1.10 1.21 1.41 - 431.000008 105 9 289 599 942 494 154 - Valence Electron Configuration # Valence Electrons Bond Length Bond Order BondEnergy(kJ/mole) Mole-cule H2H2H2H2 (  2s ) 2 (  2s *) 2 (  2p z ) 2 (  2p ) 4 (  2p *) 4 (  2s ) 2 (  2s *) 2 (  2p z ) 2 (  2p ) 4 (  2p *) 2 (  2s ) 2 (  2s *) 2 (  2p ) 4 (  2p z ) 2 (  2s ) 2 (  2s *) 2 (  2p ) 4 (  2s ) 2 (  2s *) 2 (  2p ) 2 (  2s ) 2 (  2s *) 2 (  2s ) 2 (  1s ) 2 (  1s *) 2 (  1s ) 2 (  2s ) 2 (  2s *) 2 (  2p z ) 2 (  2p ) 4 (  2p *) 4 (  2p z *) 2          

6. Bonding in Polyatomic Molecules Basically two ways to approach polyatomics. First is to use delocalized M.O.’s where e - are not confined to a single bond (region between 2 atoms) but can wander over 3 or more atoms. We will use this approach later for C bonding. Second is to use hybridization of atomic orbitals and then use these to form localized (usually) bonds.

7. Localized BeH2 orbitals: + Be + + + sp + sp - Add sp+ + 1s to get localized Be-H bond } Add sp- + 1s to get localized Be-H bond } HH.... – – + +

8. H 1s + (sp- )H 1s + (sp+ ) BeHHBe and +.. +.. ––

9. BH3 Fragment: sp 2 hybrid - + 120˚ Boron Nucleus Second sp 2 points along this direction Third sp 2 points along this direction

10. 3 sp2 hybrids 120˚

11. Overlap with H 1s to give 3 B-H bonds in a plane pointing at 120º with respect to each other: BH3 B Note B has 3 valence electrons 1s22s22p + + ++H +H+H 120˚

12. 2s + 2p x + 2p y + 2p z  sp 3 gives 4 hybrid orbitals which point to the corners of a tetrahedron. Angle between is 109º28’ sp 3 [1/4 s, 3/4 p]. Tetrahedral hybrids. 4 H atoms have 4 x 1s  4 valence e -

13. E 2s 2 2p 1s 2 sp 3 1s 2

14. Geometry of carbon sp3/H 1s Bonds in methane (CH4): sp3 hybridization on C leads to 4 bonds. CH4 is a good example. C H H H H

15. Summary of Hybridization Results BeH22 splinear H-Be-H BH33 sp2trig. plane (120º H-B-H angle) CH44 sp3tetrahedral (109º28’ H-C-H angles) ExampleGroups Attached HybridGeometry to Center Atom

16. 2p z 2p y 2p x 1s Localized Bonds and Lone Pair Electrons NH 3 This predicts 90º geometry 3H, 1s (no choice)N: 1s 2 2s 2 2p 3 Nitrogen nucleus H atom 1s orbital

17. N 107˚ H H H Geometry of NH3 found to be Trigonal Pyramidal

18. 4 N sp3 orbitals combine with 3 1s H orbitals to give 3 sp3, 1s M.O.’s leaving one sp3 hybrid left Of 5 valence e - in N, 2 go into one sp3 orbital, 3 go into other 3 sp3,s. (combined with H (1s)) One of the driving forces for the tetrahedral configuration is that it puts bonding and lone pair electron groups as far away from each other as possible.

19. N H H H Lone Pair

20. N Electron pair repulsion effect is largest for small central atoms like B, N, O. As go to larger central atoms (e.g. S or metals) frequently find this effect not so large and start to get things closer to pure p orbital bonds (90º structures) For example, H2S has H-S-H angle of 92º.

21. O H H Lone pairs stick into and out of paper H2O: 4 sp3 hybrids. Use 2 to make MO bonds with H 1s. Predicts H-O-H bond of 109º28” - Find 105º O 1s 2 2s 2 2p 4 (6 valence e - )2H 1s (2e - )