Chemistry 100 Chapter 9

1. Chemistry 100 Chapter 9 Molecular Geometry and Bonding Theories

2. Molecular Geometry The three-dimensional arrangement of atoms in a molecule  molecular geometry Lewis structures can’t be used to predict geometry Repulsion between electron pairs (both bonding and non-bonding) helps account for the molecular structure!

3. The VSEPR Model • Electrons are negatively charged, they want to occupy positions such that electron • Electron interactions are minimized as much as possible • Valence Shell Electron-Pair Repulsion Model • treat double and triple bonds as single domains • resonance structure - apply VSEPR to any of them • formal charges are usually omitted

4. Four Electron Domains – Three Different Geometries • Replacement of bonding domains (B) with nonbonding domains (E)results in a different molecular geometry. AB4 AB3E AB2E2

5. Molecules With More Than One Central Atom • Carbon #1 – tetrahedral • Carbon #2 – trigonal planar We simply apply VSEPR to each ‘central atom’ in the molecule.

6. Dipole Moments +H-F The HF molecule has a bond dipole – a charge separation due to the electronegativity difference between F and H. The shape of a molecule and the magnitude of the bond dipole(s) can give the molecule an overall degree of polarity ® dipole moment.

7. Homonuclear diatomics ® no dipole moment (O2, F2, Cl2, etc) Triatomic molecules (and greater). Must look at the net effect of all the bond dipoles. In molecules like CCl4 (tetrahedral) BF3 (trigonal planar) all the individual bond dipoles cancel Þ no resultant dipole moment.

8. Bond Dipoles in Molecules

9. More Bond Dipoles

10. Valence Bond Theory and Hybridisation • Valence bond theory • description of the covalent bonding and structure in molecules. • Electrons in a molecule occupy the atomic orbitals of individual atoms. • The covalent bond results from the overlap of the atomic orbitals on the individual atoms

11. The Bonding in Diatomic Molecules • Hydrogen molecule • a single bond between the two H 1s orbitals • a  bond • Hydrogen Chloride • a single  bond from the overlap of the Cl 3p orbital with the H 1s orbital • Chlorine molecule • a single  bond from the overlap of the Cl 3p orbitals

12. Hybrid Atomic Orbitals • Bonding and geometry in polyatomic molecules may be explained in terms of the formation of hybrid atomic orbitals • Bonds  overlap of the hybrid atomic orbitals on central atoms with appropriate half-filled atomic orbital on the terminal atoms. Look at the bonding picture in methane (CH4).

13. The CH4 Molecule

14. The Formation of the sp3 Hybrids • Mix 3 “pure” p orbitals and a “pure” s orbital • form an sp3 “hybrid” orbital. • Rationalize the bonding around the C central atom.

15. sp2 Hybridisation • Examine BH3 • (a trigonal planar molecule)

16. sp Hybridisation Examine BeF2 (a linear molecule). These sp hybrid orbitals have an angle of 180 between them.

17. A Linear Molecule The BeF2molecule

18. Double Bonds Look at ethene C2H4. Each central atom is an AB3 system, the bonding picture must be consistent with VSEPR theory.

19. Sigma () Bonds • Sigma bonds are characterized by • Head-to-head overlap. • Cylindrical symmetry of electron density about the internuclear axis.

20. Pi () Bonds • Pi bonds are characterized by • Side-to-side overlap. • Electron density above and below the internuclear axis.

21. Bond overlaps in C2H4 There are three different types of bonds [sp2 (C ) – 1s (H) ] x 4  type [sp2 (C 1 ) – sp2 (C 2 ) ]  type [2pz(C 1 ) – 2pz(C 2 ) ] p type

22. The C2H4 Molecule

23. The Bond Angles in C2H4 Any double bond  one  bond and a p bond • Bond angles HCH = HCC  120. • p bond is perpendicular to the plane containing the molecule. • Double bonds – • Rationalize by assuming sp2 hybridization exists on the central atoms!

24. The Triple Bond in C2H2 Triple bond rationalized by assuming sp hybridization exists on the central atoms! • Bond angles HCH = HCC = 180. • p bonds are perpendicular to the molecular plane. • Triple bond  one  bond and two p bonds

25. Bond Overlaps in C2H2 • There are again three different types of bonds [sp (C ) – 1s (H) ] x 2  type [sp(C 1 ) – sp(C 2 ) ]  type [2py(C 1 ) – 2py(C 2 ) ] p type [2pz(C 1 ) – 2pz(C 2 ) ] p type

26. Bonding in H2O Bonding Overlaps [sp3(O)–1s(H)] x 2  

27. Bond Overlaps in H2CO There are again three different types of bonds [sp(C) – 1s (H) ] x 2  type [sp2 (C) – sp2 (O) ]  type [2p(C) – 2p(O) ] p type

28. Key Connection – VSEPR and Valence Bond Theory!!

29. sp3d Hybridisation How can we use the hybridisation concept to explain the bonding picture PCl5. There are five bonds between P and Cl (all s type bonds). 5 sp3d orbitals® these orbitals overlap with the 3p orbitals in Cl to form the 5 s bonds with the required VSEPR geometry ® trigonal bipyramid. Bond overlaps [sp3d(P ) – 3pz (Cl) ] x 5  type

30. sp3d2 Hybridisation Look at the SF6 molecule. 6 sp3d2 orbitals® these orbitals overlap with the 2pz orbitals in F to form the 6 s bonds with the required VSEPR geometry ® octahedral. Bond overlaps [sp3d2 (S ) – 2pz (F) ] x 6  type

31. Notes for Understanding Hybridization Applied to atoms in molecules only Number hybrid orbitals = number of atomic orbitals used to make them Hybrid orbitals have different energies and shapes from the atomic orbitals from which they were made. Hybridization requires energy for the promotion of the electron and the mixing of the orbitals ® energy is offset by bond formation.

32. Delocalised Bonding • Valence bond theory – • bonding electrons have been totally associated with the two atoms that form the bond  they are localized. • What about the bonding situation in benzene, the nitrate ion, the carbonate ion?

33. Bonding in Aromatic Molecules • Benzene • C-C s bonds are formed from the sp2 hybrid orbitals. • Unhybridized 2pz orbital on adjacent C atoms overlap (bonds).

34. Bonding in the Benzene Molecule • The p bonds extend over the whole molecule • the p electronsbonds are delocalized – they are free to move around the benzene ring. • Resonance structures – delocalization of the -electrons.

35. The Nitrate Anion • Three resonance structures • Alternating single and double bonds • Blend resonance structures • Delocalized  bond over anion backbone

36. Molecular Orbital (M.O.) Theory • Valence bond and the concept of the hybridisation of atomic orbitals does not account for a number of fundamental observations of chemistry. • To reconcile these and other differences, we turn to molecular orbital theory (MO theory). • MO theory – covalent bonding is described in terms of molecular orbitals • the combination of atomic orbitals that results in an orbital associated with the whole molecule.

37. Constructive and Destructive Interference Constructive + Destructive +

38. ybonding = C1 ls (H 1) + C2 ls (H 2) yanti = C1 ls (H 1) - C2 ls (H 2) Bonding Orbital® a centro-symmetric orbital (i.e. symmetric about the line of symmetry of the bonding atoms). Bonding M’s have lower energy and greater stability than the AO’s from which it was formed. Electron density is concentrated in the region immediately between the bonding nuclei.

39. Anti-bonding orbital® a node (0 electron density) between the two nuclei. In an anti-bonding MO, we have higher energy and less stability than the atomic orbitals from which it was formed. As with valance bond theory (hybridisation) 2 AO’s ® 2 MO’s

40. Bonding and Anti-Bonding M.O.’s from 1s atomic Orbitals

41. The MO’s in the H2 Atom

42. The situation for two 2s orbitals is the same! The situation for two 3s orbital is the same. • Let’s look at the following series of molecules H2, He2+, He2 bond order = ½ {bonding - anti-bonding e-‘s}. • Higher bond order º greater bond stability.