1. Polarity and IMF

2. Polar Bonds When the atoms in a bond are the same, the electrons are shared equally. This is a nonpolar covalent bond. When two different atoms are connected, the atoms may not be shared equally. This is a polar covalent bond. How do we measure how strong the atoms pull on electrons?

3. Electronegativity A measure of how strongly the atoms attract electrons in a bond. The bigger the electronegativity difference the more polar the bond. 0.0 - 0.3 Covalent nonpolar 0.3 - 1.67 Covalent polar >1.67 Ionic

4. How to show a bond is polar Isn’t a whole charge just a partial charge  means a partially positive  means a partially negative The Cl pulls harder on the electrons The electrons spend more time near the Cl H Cl  

5. Polar Molecules Molecules with ends

6. Polar Molecules Molecules with a positive and a negative end Requires two things to be true ¬ The molecule must contain polar bonds This can be determined from differences in electronegativity. ­ Symmetry can not cancel out the effects of the polar bonds. Must determine geometry first.

7. Is it polar? HF H 2 O NH 3 CCl 4 CO 2

8. Intermolecular Forces What holds molecules to each other

9. Intermolecular Forces They are what make solid and liquid molecular compounds possible. The weakest are called van der Waal’s forces - there are two kinds Dispersion forces Dipole Interactions – depend on the number of electrons – more electrons stronger forces – Bigger molecules

10. Dipole interactions Depend on the number of electrons More electrons stronger forces Bigger molecules more electrons Fluorine is a gas Bromine is a liquid Iodine is a solid

11. Dipole interactions Occur when polar molecules are attracted to each other. Slightly stronger than dispersion forces. Opposites attract but not completely hooked like in ionic solids.

12. Dipole interactions Occur when polar molecules are attracted to each other. Slightly stronger than dispersion forces. Opposites attract but not completely hooked like in ionic solids. HFHF  HFHF 

13. Dipole Interactions    

14. Hydrogen bonding Are the attractive force caused by hydrogen bonded to F, O, or N. F, O, and N are very electronegative so it is a very strong dipole. The hydrogen partially share with the lone pair in the molecule next to it. The strongest of the intermolecular forces.

15. Hydrogen Bonding H H O ++ -- ++ H H O ++ -- ++

16. Hydrogen bonding H H O H H O H H O H H O H H O H H O H H O

17. MOLECULAR SHAPES OF COVALENT COMPOUNDS

18. VSepR tHEORY

19. What Vsepr means Since electrons do not like each other, because of their negative charges, they orient themselves as far apart as possible, from each other. This leads to molecules having specific shapes.

20. Things to remember Atoms bond to form an Octet (8 outer electrons/full outer energy level) Bonded electrons take up less space then un-bonded/unshared pairs of electrons.

22. Linear Number of Bonds = 2 Number of Shared Pairs of Electrons = 2 Bond Angle = 180° EXAMPLE: BeF 2

23. Trigonal Planar Number of Bonds = 3 Number of Shared Pairs of Electrons = 3 Number of Unshared Pairs of Electrons = 0 Bond Angle = 120° EXAMPLE: GaF 3

24. Tetrahedral Number of Bonds = 4 Number of Shared Pairs of Electrons = 4 Number of Unshared Pairs of Electrons = 0 Bond Angle = 109.5° EXAMPLE: CH 4

25. Trigonal Pyramidal Number of Bonds = 3 Number of Shared Pairs of Electrons = 4 Number of Unshared Pairs of Electrons = 1 Bond Angle = <109.5° EXAMPLE: NH 3

26. Bent #1 Number of Bonds = 2 Number of Shared Pairs of Electrons = 2 Number of Unshared Pairs of Electrons = 2 Bond Angle = < 120° EXAMPLE: H 2 O