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CH2. Molecules and covalent bonding Lewis Structures VSEPR MO Theory 1 Lewis structure H3PO4 • Skeleton is: • Count total valence electrons: 1P= 5 3H= 3 4 O = 24 Total = 32 e- or 16 valence e- pairs. • 7 e- pairs needed to form s skeleton. 2 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. 3 Resonance in NO3- experimental data - nitrate is planar with 3 equivalent N-O bonds 4 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 3 4 5 6 7 8 linear equilateral triangular tetrahedral (Td) trigonal bipyramidal (TBP) octahedral (Oh) pentagonal bipyramidal square antiprismatic HF2BF3 CF4 PF5 SF6 IF7 TaF835 Drawing Oh and Td molecules It's often useful to draw octahedra and tetrahedra with a cubic framework 6 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) 7 POCl3 note that in :PCl3 the <ClPCl = 100.3, the lone pair is more repulsive towards other ligands than the multiple bond ! based on Td geometry < ClPCl = 103.3° due to repulsion by multiple bond Ligands move away from multiple bond 8 XeF5+ 5 Xe-F bonds and 1 lone pair on Xe geometry based on Oh coordination lone pair repulsion gives < FeqXeFeq = 87° < FaxXeFeq = 78° 9 Fajan’s rule bond polarization is towards ligands with higher c, decreasing repulsive effect. Lone pairs are the most repulsive. ex: NH3 vs NF3 < HNH = 107.3° < FNF = 102.1° 10 Inert pair effect • In Sn Sb Te Tl Pb Bi • 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. 11 Inert pair effect - evidence • Bond angles near 90°: NH3 107.2 AsH3 91.8 SbH3 91.3 • PbO unit cell • H2O H2Se H2Te 104.5 91 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 SnF2) ex: In and Tl both form monochlorides, B, Al, Ga form trichlorides. Vacant coordination sites where the lone pair resides ex: PbO 12 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 13 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) 14 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. 15 H2 16 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. 17 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 nonbonding. 18 I and Ea for atoms and diatomics species I (kJ/mol) N O O2 NO F F2 C C2 1402 1314 1165 893 1681 1515 1086 Ea 142 43 123 300 19 Li2 – F2 MO’s 20 Some diatomic bond data bond order r0 in pm D0 in kJ/mol O2 2 121 494 O2 - 1½ 126 O22- 1 149 F2 O2+ NO NO+ N2 1 2½ 2½ 3 3 142 112 115 106 110 155 942 21 Spectroscopic data for MO’s 22 HF 23 Ketalaar triangle HF 24 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. 25 H3+ 26 BeH2 27 Correlation diagram for MH2 M < HMH Be 180° B 131 C 136 N 103 O 105 28 Bonding MO’s in H2O 29 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. 30 NH3 --calculated MO diagram 31 SF6 See textbook Resource Section 5 for SALCs 32