1. 1 CH2. Molecules and covalent bonding Lewis Structures VSEPR MO Theory

2. 2 Lewis structure H 3 PO 4 Skeleton is: Count total valence electrons: 1 P = 5 3 H = 3 4 O = 24 Total = 32 e - or 16 valence e - pairs. 7 e - pairs needed to form  skeleton.

3. 3 Lewis structure H 3 PO 4 Add remaining e - pairs: Left has a formal charge of +1 on P and -1 on one O, right has 5 e - pairs around P (hypervalence) Analysis of phosphoric acid shows purely Td phosphate groups, which requires something beyond either simple Lewis model.

4. 4 Resonance in NO 3 - experimental data - nitrate is planar with 3 equivalent N-O bonds

5. 5 VSEPR model Count e - pairs about the central atom (draw Lewis structure if needed). Include non-bonding pairs, but not multiple bonds. Geometry maximizes separation: # e pairsgeometry example 2linear HF 2 - 3equilateral triangular BF 3 4tetrahedral (Td) CF 4 5trigonal bipyramidal (TBP) PF 5 6octahedral (Oh) SF 6 7pentagonal bipyramidal IF 7 8square antiprismatic TaF 8 3-

6. 6 Drawing Oh and Td molecules It's often useful to draw octahedra and tetrahedra with a cubic framework

7. 7 Deviations from ideal geometries: unshared pairs and multiple bonds require larger bite ex: CH 4, NH 3, H 2 O <H-C-H = 109.5°, <H-N-H = 107.3, <H-O-H = 104.5 ex: ICl 4 - 6 e pairs around I, 2 lone pairs and 4 e pair bonds to Cl Oh coordination, and geometry is square planar (lone pairs are trans, not cis)

8. 8 POCl 3 based on Td geometry < ClPCl = 103.3° due to repulsion by multiple bond note that in :PCl 3 the <ClPCl = 100.3, the lone pair is more repulsive towards other ligands than the multiple bond ! Ligands move away from multiple bond

9. 9 XeF 5 + 5 Xe-F bonds and 1 lone pair on Xe geometry based on Oh coordination lone pair repulsion gives < F eq XeF eq = 87° < F ax XeF eq = 78°

10. 10 Fajan’s rule bond polarization is towards ligands with higher , decreasing repulsive effect. Lone pairs are the most repulsive. ex: NH 3 vs NF 3 < HNH = 107.3° < FNF = 102.1°

11. 11 Inert pair effect VSEPR geometries require hybridization (valence bond term) or linear combinations (MO term) of central atom orbitals. For example, Td angles require sp 3 hybrid orbitals. More on this in MO theory section. Period 5 and 6 p-block central atoms often show little hybridization (ex: they form bond with orbitals oriented at 90° as in purely p orbitals). This can be ascribed to the weaker bonding of larger atoms to ligands. In Sn Sb Te Tl Pb Bi

12. 12 Inert pair effect - evidence Bond angles near 90°: NH 3 107.2 H 2 O 104.5 AsH 3 91.8 H 2 Se 91 SbH 3 91.3 H 2 Te 89.5 Increased stability of lower oxidation states ex: Si, and Ge are generally 4+, but Sn and Pb are common as 2+ ions (as in stannous fluoride SnF 2 ) ex: In and Tl both form monochlorides, B, Al, Ga form trichlorides. Vacant coordination sites where the lone pair resides ex: PbO PbO unit cell

13. 13 Fluxionality PF 5 if TBP has 2 types of F ligands (equatorial and axial). 19 F NMR spectra at RT show only a single peak (slightly broadened). PF 5 is fluxional at RT, i.e. the F ligands exchange rapidly, only a single "average" F ligand is seen by NMR. Only occurs if ligand exchange is faster than the analytical method. IR and Raman have shorter interaction times and show 2 types of P-F bonding at RT. Even low temp NMR studies cannot resolve two F environments

14. 14 Berry pseudo-rotation Sequences of the MD-Simulation of PF5 at 750K ( Daul, C., et al, Non-empirical dynamical DFT calculation of the Berry pseudorotation of PF5, Chem. Phys. Lett. 1996, 262, 74 )

15. 15 Molecular Orbitals Use linear combinations of atomic orbitals to derive symmetry-adapted linear combinations (SALCs). Use symmetry to determine orbital interactions. Provide a qualitative MO diagram for simple molecules. Read and analyze an MO diagram by sketching MO’s / LCAO’s, describing the geometric affect on relative MO energies.

16. 16 H2H2

17. 17 Some rules The number of AO’s and MO’s must be equal. This follows from the mathematics of independent linear combinations. More on symmetry labels later, but they come from the irreducible representations for the point group.  MO’s are symmetric about bond axis,  MO’s are not. Subscipt g is gerade (has center of symmetry), u is ungerade. Antibonding orbitals are often given a * superscript. The bond order = ½ (bonding e - - antibonding e - ). The bond energy actually depends on the energies of the filled MO’s relative to filled AO’s.

18. 18 O2O2 MO theory predicts 2 unpaired e -, this is confirmed by experiment. Bond order = ½ (8-4) = 2, as in Lewis structure. MO indicates distribution and relative energies of the MO's, Lewis structure says only bonding or non- bonding.

19. 19 I and E a for atoms and diatomics speciesI (kJ/mol)EaEa N1402 O1314142 O2O2 1165 43 NO 893 F1681 F2F2 1515 C1086123 C2C2 300

20. 20 Li 2 – F 2 MO’s

21. 21 Some diatomic bond data bond orderr 0 in pmD 0 in kJ/mol O2O2 2121494 O2-O2- 1 ½126 O 2 2- 1149 F2F2 1142155 O2+O2+ 2 ½112 NO2 ½115 NO + 3106 N2N2 3110942

22. 22 Spectroscopic data for MO’s

23. 23 HF

24. 24 Ketalaar triangle HF

25. 25 Hybridization Linear combinations of AO’s from same atom makes hybrid orbitals. Hybridization can be included in the MO diagram. In MO theory, any proportion of s and p can be mixed (the coefficients of the AO’s are variable). sp and sp 3 hybrids are specific examples.

26. 26 H3+H3+

27. 27 BeH 2

28. 28 Correlation diagram for MH 2 M < HMH Be 180° B 131 C 136 N 103 O 105

29. 29 Bonding MO’s in H 2 O

30. 30 NH 3 Use triangular H 3 MO’s from above as SALC's of the H ligand orbitals. Must relabel to conform with lower symmetry pt group C 3v. They become a 1 and e. Combine with N valence orbitals with same symmetry.

31. NH 3 --calculated MO diagram 31

32. 32 SF 6 See textbook Resource Section 5 for SALCs