CH2. Molecules and covalent bonding Lewis Structures VSEPR MO Theory

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

2. Lewis structure H3PO4 • 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 s skeleton.

3. Lewis structure H3PO4 • 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. Resonance in NO3- • experimental data - nitrate is planar with 3 equivalent N-O bonds

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 pairs geometry example 2 linear HF2- 3 equilateral triangular BF3 4 tetrahedral (Td) CF4 5 trigonal bipyramidal (TBP) PF5 6 octahedral (Oh) SF6 7 pentagonal bipyramidal IF7 8 square antiprismatic TaF83-

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

7. Deviations from ideal geometries: unshared pairs and multiple bonds require larger bite ex: CH4, NH3, H2O <H-C-H = 109.5°, <H-N-H = 107.3, <H-O-H = 104.5 ex: ICl4- 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. Ligands move away from multiple bond POCl3 based on Td geometry < ClPCl = 103.3° due to repulsion by multiple bond note that in :PCl3 the <ClPCl = 100.3, the lone pair is more repulsive towards other ligands than the multiple bond !

9. XeF5+ 5 Xe-F bonds and 1 lone pair on Xe geometry based on Oh coordination lone pair repulsion gives < FeqXeFeq = 87° < FaxXeFeq = 78°

10. Fajan’s rule bond polarization is towards ligands with higher c, decreasing repulsive effect. Lone pairs are the most repulsive. ex: NH3vs NF3 < HNH = 107.3° < FNF = 102.1°

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 sp3 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. Inert pair effect - evidence • Bond angles near 90°: NH3     107.2           H2O     104.5 AsH3     91.8           H2Se      91 SbH3     91.3           H2Te     89.5 • Increased stability of lower oxidation statesex: Si, and Ge are generally 4+, but Sn and Pb are common as 2+ ions (as in stannous fluoride SnF2) 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. Fluxionality • PF5 if TBP has 2 types of F ligands (equatorial and axial). • 19F NMR spectra at RT show only a single peak (slightly broadened). • PF5 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. 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. 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. H2

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. s MO’s are symmetric about bond axis, p 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. O2 • 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. species I (kJ/mol) Ea N 1402 O 1314 142 O2 1165 43 NO 893 F 1681 F2 1515 C 1086 123 C2 300 I and Ea for atoms and diatomics

20. Li2 – F2 MO’s

21. bond order r0 in pm D0 in kJ/mol O2 2 121 494 O2- 1 ½ 126 O22- 1 149 F2 1 142 155 O2+ 2 ½ 112 NO 2 ½ 115 NO+ 3 106 N2 3 110 942 Some diatomic bond data

22. Spectroscopic data for MO’s

23. HF

24. Ketalaar triangle HF

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 sp3 hybrids are specific examples.

26. H3+

27. BeH2

28. Correlation diagram for MH2 M < HMH Be 180° B 131 C 136 N 103 O 105

29. Bonding MO’s in H2O

30. NH3 Use triangular H3 MO’s from above as SALC's of the H ligand orbitals. Must relabel to conform with lower symmetry pt group C3v. They become a1 and e. Combine with N valence orbitals with same symmetry.

31. NH3 --calculated MO diagram

32. SF6 See textbook Resource Section 5 for SALCs