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

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

3.  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 - n Exceptions: A. Octet Rule F B F F H O HN O Very unstable!! F F S F F

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

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

6. B. Drawing Lewis Diagrams n BeCl 2 1 Be × 2e - = 2e - 2 Cl × 7e - = 14e - 16e - Cl Be Cl - 4e - 12e -

7. B. Drawing Lewis Diagrams n CO 2 1 C × 4e - = 4e - 2 O × 6e - = 12e - 16e - O C O - 4e - 12e -

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

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

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

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

12. D. Resonance Structures O O S O O O S O O O S O n SO 3

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

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

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

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

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

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

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

20. 3 total 2 bond 1 lone BENT <120° SO 2

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

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

23. 4 total 2 bond 2 lone BENT 104.5° H2OH2O C. Common Molecular Shapes

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

25. n CO 2 O C O 2 total 2 bond 0 lone LINEAR 180° D. Examples

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

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

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

29. B. Determining Molecular Polarity n Nonpolar Molecules  Dipole moments are symmetrical and cancel out. BF 3 F F F B

30. B. Determining Molecular Polarity n Polar Molecules  Dipole moments are asymmetrical and don’t cancel. net dipole moment H2OH2O H H O

31. CHCl 3 H Cl B. Determining Molecular Polarity n Therefore, polar molecules have...  asymmetrical shape (lone pairs) or  asymmetrical atoms net dipole moment

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

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

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