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 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. AB 4 AB 3 EAB 2 E 2

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

6. Dipole Moments  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.  + H-F 

7.  Homonuclear diatomics  no dipole moment (O 2, F 2, Cl 2, etc)  Triatomic molecules (and greater). Must look at the net effect of all the bond dipoles.  In molecules like CCl 4 (tetrahedral) BF 3 (trigonal planar) all the individual bond dipoles cancel  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 atomscovalent bond

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  Look at the bonding picture in methane (CH 4 ). 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.

13. The CH 4 Molecule

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

15. sp 2 Hybridisation  Examine BH 3 (a trigonal planar molecule)

16. sp Hybridisation  Examine BeF 2 (a linear molecule).  These sp hybrid orbitals have an angle of 180 between them.

17. A Linear Molecule The BeF 2 molecule

18. Double Bonds  Look at ethene C 2 H 4.  Each central atom is an AB 3 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 C 2 H 4  There are three different types of bonds [sp 2 (C ) – 1s (H) ] x 4 type [sp 2 (C 1 ) – sp 2 (C 2 ) ]  type [2p z (C 1 ) – 2p z (C 2 ) ]  type

22. The C 2 H 4 Molecule

23. The Bond Angles in C 2 H 4  Bond angles HCH = HCC  120.  bond is perpendicular to the plane containing the molecule.  Double bonds – Rationalize by assuming sp 2 hybridization exists on the central atoms! Any double bond  one  bond and a  bond

24. The Triple Bond in C 2 H 2  Bond angles HCH = HCC = 180.  bonds are perpendicular to the molecular plane.  Triple bond  one  bond and two  bonds Triple bond rationalized by assuming sp hybridization exists on the central atoms!

25. Bond Overlaps in C 2 H 2  There are again three different types of bonds [sp (C ) – 1s (H) ] x 2 type [sp (C 1 ) – sp (C 2 ) ]  type [2p y (C 1 ) – 2p y (C 2 ) ]  type [2p z (C 1 ) – 2p z (C 2 ) ]  type

26. Bonding in H 2 O Bonding Overlaps [sp 3 (O)–1s(H)] x 2  

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

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

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

30. sp 3 d 2 Hybridisation  Look at the SF 6 molecule.  6 sp 3 d 2 orbitals  these orbitals overlap with the 2p z orbitals in F to form the 6  bonds with the required VSEPR geometry  octahedral.  Bond overlaps [sp 3 d 2 (S ) – 2p z (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  bonds are formed from the sp 2 hybrid orbitals. Unhybridized 2p z orbital on adjacent C atoms overlap (bonds).

34. Bonding in the Benzene Molecule  The bonds extend over the whole molecule the electrons bonds 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.  bonding = C 1  ls (H 1) + C 2 ls (H 2)  anti = C 1  ls (H 1) - C 2 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 H 2 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 H 2, He 2 +, He 2 bond order = ½ {bonding - anti-bonding e - ‘s}.  Higher bond order  greater bond stability.