I. Lewis Diagrams (p. 184-189)

1. Ch. 6 – Molecular Structure I. Lewis Diagrams(p. 184-189)

2. A. Octet Rule • Remember… • Most atoms form bonds in order to have 8 valence electrons.

3. F F F S F F F F B F F H O H N O Very unstable!! A. Octet Rule • Exceptions: • Hydrogen  2 valence e- • Groups 1,2,3 get 2,4,6 valence e- • Expanded octet  more than 8 valence e- (e.g. S, P, Xe) • Radicals  odd # of valence e-

4. B. Drawing Lewis Diagrams • Find total # of valence e-. • Arrange atoms - singular atom is usually in the middle. • Form bonds between atoms (2 e-). • Distribute remaining e- to give each atom an octet (recall exceptions). • If there aren’t enough e- to go around, form double or triple bonds.

5. B. Drawing Lewis Diagrams • CF4 1 C × 4e- = 4e- 4 F × 7e- = 28e- 32e- F F C F F - 8e- 24e-

6. B. Drawing Lewis Diagrams • BeCl2 1 Be × 2e- = 2e- 2 Cl × 7e- = 14e- 16e- ClBeCl - 4e- 12e-

7. B. Drawing Lewis Diagrams • CO2 1 C × 4e- = 4e- 2 O × 6e- = 12e- 16e- OCO - 4e- 12e-

8. C. Polyatomic Ions • To find total # of valence e-: • Add 1e- for each negative charge. • Subtract 1e- for each positive charge. • Place brackets around the ion and label the charge.

9. C. Polyatomic Ions • ClO4- 1 Cl × 7e- = 7e- 4 O × 6e- = 24e- 31e- O O Cl O O + 1e- 32e- - 8e- 24e-

10. C. Polyatomic Ions • NH4+ 1 N × 5e- = 5e- 4 H × 1e- = 4e- 9e- H H N H H - 1e- 8e- - 8e- 0e-

11. D. Resonance Structures • Molecules that can’t be correctly represented by a single Lewis diagram. • Actual structure is an average of all the possibilities. • Show possible structures separated by a double-headed arrow.

12. O O S O O O S O O O S O D. Resonance Structures • SO3

13. Ch. 6 – Molecular Structure II. Molecular Geometry(p. 197-200)

14. A. VSEPR Theory • Valence Shell Electron Pair Repulsion Theory • Electron pairs orient themselves in order to minimize repulsive forces.

15. Lone pairs repel more strongly than bonding pairs!!! A. VSEPR Theory • Types of e- Pairs • Bonding pairs - form bonds • Lone pairs - nonbonding e-

16. Bond Angle A. VSEPR Theory • Lone pairs reduce the bond angle between atoms.

17. Know the 8 common shapes & their bond angles! B. Determining Molecular Shape • Draw the Lewis Diagram. • Tally up e- pairs on central atom. • double/triple bonds = ONE pair • Shape is determined by the # of bonding pairs and lone pairs.

18. BeH2 C. Common Molecular Shapes 2 total 2 bond 0 lone LINEAR 180°

19. BF3 C. Common Molecular Shapes 3 total 3 bond 0 lone TRIGONAL PLANAR 120°

20. SO2 C. Common Molecular Shapes 3 total 2 bond 1 lone BENT <120°

21. CH4 C. Common Molecular Shapes 4 total 4 bond 0 lone TETRAHEDRAL 109.5°

22. NH3 C. Common Molecular Shapes 4 total 3 bond 1 lone TRIGONAL PYRAMIDAL 107°

23. H2O C. Common Molecular Shapes 4 total 2 bond 2 lone BENT 104.5°

24. F P F F D. Examples • PF3 4 total 3 bond 1 lone TRIGONAL PYRAMIDAL 107°

25. OCO D. Examples • CO2 2 total 2 bond 0 lone LINEAR 180°

26. Ch. 6 – Molecular Structure III. Polarity & IMF(p. 204-207)

27. + - H Cl A. Dipole Moment • Direction of the polar bond in a molecule. • Arrow points toward the more e-neg atom.

28. B. Determining Molecular Polarity • Depends on: • dipole moments • molecular shape

29. F BF3 B F F B. Determining Molecular Polarity • Nonpolar Molecules • Dipole moments are symmetrical and cancel out.

30. O net dipole moment H2O H H B. Determining Molecular Polarity • Polar Molecules • Dipole moments are asymmetrical and don’t cancel .

31. H net dipole moment CHCl3 Cl Cl Cl B. Determining Molecular Polarity • Therefore, polar molecules have... • asymmetrical shape (lone pairs) or • asymmetrical atoms

32. Dipole-Dipole Forces • Attractive forces between polar covalent molecules

33. London (Dispersion) Forces • Attractive forces between the electron clouds of large molecules in large quantity • Larger mass = Larger London Forces

34. Hydrogen Bonding • Special dipole-dipole attraction that involves H bonded with high electronegative elements N, O, or F