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 minimised 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. 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.

5. 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.

6. 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.

7. Bond Dipoles in Molecules

8. More Bond Dipoles

9. 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

10. The Bonding in H2 1s (H1) – 1s(H2)  bond Overlap Region • Hydrogen molecule • a single bond indicating the overlap of the 1s orbitals on the individual atoms • cylindrical symmetry with respect to the line joining the atomic centres, i.e., a  bond

11. The Bonding in H2

12. The Cl2 Molecule Bonding description 3pz (Cl 1) – 3pz (Cl 2) In the chlorine molecule, we observe a single bond indicating the overlap of the 3p orbitals on the individual atoms.

13. Is This a s Bond?

14. 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).

15. The CH4 Molecule

16. 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.

17. sp2 Hybridisation We mix 2 “pure” p orbitals and a “pure” s orbital to form “hybrid” or mixed sp2 orbitals. These three sp2 hybrid orbitals lie in the same plane with an angle of 120 between them. What if we try to rationalise the bonding picture in the BH3 (a trigonal planar molecule)?

18. A Trigonal Planar Molecule Overlap region Overlap regions

19. sp Hybridisation What if we try to rationalize the bonding picture in the BeH2 species (a linear molecule)? We mix a single “pure” p orbital and a “pure” s orbital to form two “hybrid” or mixed sp orbitals These sp hybrid orbitals have an angle of 180 between them.

20. A Linear Molecule Overlap Regions The BeH2 molecule

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

22. The  Bond • Additional feature • an unhybridized p orbital on adjacent carbon atoms. • Overlap the two parallel 2pz orbitals (a p-bond is formed).

23. 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

24. The C2H4 Molecule

25. 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!

26. 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

27. 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

28. 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

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

30. 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

31. 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

32. Notes for Understanding Hybridisation • 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. • Hybridisation requires energy for the promotion of the electron and the mixing of the orbitals ® energy is offset by bond formation.

33. Delocalised Bonding In almost all the cases where we described the bonding n the molecule, the bonding electrons have been totally associated with the two atoms that form the bond  they are localised. What about the bonding situation in benzene, the nitrate ion, the carbonate ion? In benzene, the C-C s bonds are formed from the sp2 hybrid orbitals. The unhybridised 2pz orbital on C overlaps with another 2pz orbital on the adjacent C atom.

34. Three p bonds are formed. These p bonds extend over the whole molecule (i.e. the p bonds are delocalised). The p electrons are free to move around the benzene ring. Any species where we had several resonance structures, we would have delocalisation of the -electrons.

35. Delocalised Electrons in Molecules

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 interference  the two e- waves interact favourably; loosely analogous to a build-up of e- density between the two atomic centres. destructive interference unfavourable interaction of e- waves; analogous to the decrease of e- density between two atomic centres. Recall the wave properties of electrons.

38. Constructive and Destructive Interference Constructive + Destructive +

39. 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.

40. 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

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

42. The MO’s in the H2 Atom

43. 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.