1. IIIIII Molecular Geometry (p. 232 – 236) Ch. 8 – Molecular Structure

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

3. A. VSEPR Theory  Types of e - Pairs  Bonding pairs – form bonds  Lone pairs – nonbonding e -  Total e - pairs– bonding + lone pairs Lone pairs repel more strongly than bonding pairs!!!

4. A. VSEPR Theory  Lone pairs reduce the bond angle between atoms Bond Angle

5.  Draw the Lewis Diagram  Tally up e - pairs on central atom (bonds + lone pairs)  double/triple bonds = ONE pair  Shape is determined by the # of bonding pairs and lone pairs Know the 13 common shapes & their bond angles! B. Determining Molecular Shape

6. C. Common Molecular Shapes # 1 2 total 2 bond 0 lone LINEAR 180° BeH 2 → Electronic Geometry = linear Hybridization = sp

7. 3 total 3 bond 0 lone TRIGONAL PLANAR 120° BF 3 C. Common Molecular Shapes # 2 → Electronic Geometry = trigonal planar Hybridization = sp 2

8. C. Common Molecular Shapes # 3 3 total 2 bond 1 lone BENT <120° NO 2 1- → Electronic Geometry = trigonal planar Hybridization = sp 2

9. 4 total 4 bond 0 lone TETRAHEDRAL 109.5° CH 4 C. Common Molecular Shapes # 4 → Electronic Geometry = tetrahedral Hybridization = sp 3

10. 4 total 3 bond 1 lone TRIGONAL PYRAMIDAL 107° NCl 3 C. Common Molecular Shapes # 5 → Electronic Geometry = tetrahedral Hybridization = sp 3 <109.5°

11. 4 total 2 bond 2 lone BENT 104.5° H2OH2O C. Common Molecular Shapes # 6 → Electronic Geometry = tetrahedral Hybridization = sp 3 <109.5°

12. 5 total 5 bond 0 lone TRIGONAL BIPYRAMIDAL 120°/90° PI 5 C. Common Molecular Shapes # 7 → Electronic Geometry = trigonal bipyramidal Hybridization = sp 3 d

13. 6 total 6 bond 0 lone OCTAHEDRAL 90° SH 6 C. Common Molecular Shapes # 11 → Electronic Geometry = octahedral Hybridization = sp 3 d 2

14.  SeO 3 3 total 3 bond 0 lone D. Examples O O Se O E.G. = TRIGONAL PLANAR M.G. = TRIGONAL PLANAR 120° Hybridization = sp 2

15.  AsH 3 4 total 3 bond 1 lone E.G. = TETRAHEDRAL M.G. = TRIGONAL PYRAMIDAL 107° (<109.5°) H As H H D. Examples Hybridization = sp 3

16. A. Bond Polarity n Most bonds are a blend of ionic and covalent characteristics n Difference in electronegativity determines bond type

17. A. Bond Polarity n Electronegativity  Attraction an atom has for a shared pair of electrons  higher e - neg atom   -  lower e - neg atom   +  Draw the Lewis structure for HCl & label partial charges

18. B. Molecular Polarity n Polar molecule = one end slightly + and one end slightly – n Molecule with 2 poles = dipolar molecule or dipole

19. B. Molecular Polarity n Shape, symmetry and bond polarity determines molecular polarity n H – O bond is polar and water is asymmetrical, so H 2 O is polar n C – Cl bond is polar, but CCl 4 is symmetrical, so molecule is nonpolar

20. B. Molecular Polarity n Identify each molecule as polar or nonpolar  SCl 2 O2O2  CS 2  CF 4  CH 2 F 2 Tetrahedral, bent → polar Nonpolar bonds → nonpolar Linear → nonpolar Tetrahedral → nonpolar Tetrahedral → polar

21. C. Definition of IMF n IMF = Intermolecular Forces n Attractive forces between molecules  Much weaker than chemical bonds within molecules

22. D. Types of IMF Van der Waals

23. D. Types of IMF n Van der Waals n The force of attraction between nonpolar molecules View animation online.animation

24. D. Types of IMF n Dipole-Dipole Forces n The attraction between two polar molecules + + - - View animation online.animation

25. D. Types of IMF n Hydrogen Bonding n The force of attraction that occurs when hydrogen is bound to N, O, or F an is attracted to the lone pairs on another molecule

26. n PCl 3  polar = dispersion, dipole-dipole n CH 4  nonpolar = van der Waals n HF  H-F bond = dispersion, dipole- dipole, hydrogen bonding E. Determining IMF