1. BELLRINGER Explain in complete sentences the difference
between ionic compound and polar molecule

2. REMINDER ELECTRIC DIPOLE HAS 2 CHARGES : + and –
WE CONSIDER IONIC BOND AND POLAR
MOLECULE. BOTH ARE DESCRIBED BY
ELECTRIC DIPOLE
EACH ATOM IS DESCRIBED BY
ELECTRONEGATIVITY WHICH IS – CHARGE
NEGATIVE OR LEAST ELECTRONEGATIVITY
IS POSITIVE CHARGE

3. ACTIVITY
Molecule Polarity. In this activity you will use a PhET simulation to explore molecule polarity.
 Part I: What factors affect molecule polarity?
 Explore the Molecule Polarity simulation for a few minutes. In each of the three tabs, try to find all of the controls and figure out how they work.
 Two Atoms tab. Describe all of the ways you can change the polarity of the two-atom molecule. Explain how the representations below help you understand molecule polarity.

4. Three Atoms tab. Describe any new ways you can change the polarity of the three-atom molecule. Explain the relationship between the bond dipoles and the molecular dipole. Can a non-polar molecule contain polar bonds? Explain your answer with an example.
 Real Molecules tab. Predict the polarity of 6 real molecules. First, draw the molecules and any bond dipoles. Then draw any molecular dipoles. Explain your reasoning before you check your predictions with the simulation.

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
H2O
NH3
CCl4
CO2

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 Fluorine is a gas Bromine is a liquid
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

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.

13. Dipole interactions H F d+ d- H F d+ d-
Occur when polar molecules are attracted to each other.
Slightly stronger than dispersion forces.
Opposites attract but not completely hooked like in ionic solids.
H F
d+ d-
H F
d+ d-

14. Dipole Interactions
d+ d-
d+ d-
d+ d-
d+ d-
d+ d-
d+ d-
d+ d-
d+ d-

15. 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.

16. Hydrogen Bonding H O d+ d- H O d+ d-

17. Hydrogen bonding H O H O H O H O H O H O H O

18. MOLECULAR SHAPES OF
COVALENT COMPOUNDS

19. VSepR tHEORY
ALENCE
HELL
VSEPR
LECTRON
AIR
EPULSION

20. This leads to molecules having specific shapes.
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.

21. 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. HERE ARE
THE
RESULTING
MOLECULAR
SHAPES

23. Linear BeF2 Number of Bonds = 2
EXAMPLE:
BeF2
Number of Bonds = 2
Number of Shared Pairs of Electrons = 2
Bond Angle = 180°

24. Trigonal Planar GaF3 Number of Bonds = 3
EXAMPLE:
GaF3
Number of Bonds = 3
Number of Shared Pairs of Electrons = 3
Number of Unshared Pairs of Electrons = 0
Bond Angle = 120°

25. Bent #1 H2O Number of Bonds = 2
EXAMPLE:
H2O
Number of Bonds = 2
Number of Shared Pairs of Electrons = 2
Number of Unshared Pairs of Electrons = 2
Bond Angle = < 120°

26. Bent #2 O3 Number of Bonds = 2 Number of Shared Pairs of Electrons = 2
EXAMPLE: O3
Number of Bonds = 2
Number of Shared Pairs of Electrons = 2
Number of Unshared Pairs of Electrons = 1
Bond Angle = >120°

27. Tetrahedral CH4 Number of Bonds = 4
EXAMPLE:
CH4
Number of Bonds = 4
Number of Shared Pairs of Electrons = 4
Number of Unshared Pairs of Electrons = 0
Bond Angle = 109.5°

28. Trigonal Pyramidal NH3 Number of Bonds = 3
EXAMPLE:
NH3
Number of Bonds = 3
Number of Shared Pairs of Electrons = 4
Number of Unshared Pairs of Electrons = 1
Bond Angle = <109.5°

29. NbF5 Trigonal bIPyramidal Number of Bonds = 5
EXAMPLE:
NbF5
Number of Bonds = 5
Number of Shared Pairs of Electrons = 5
Number of Unshared Pairs of Electrons = 0
Bond Angle = <120°

30. OCTAHEDRAL SF6 Number of Bonds = 6
EXAMPLE:
SF6
Number of Bonds = 6
Number of Shared Pairs of Electrons = 6
Number of Unshared Pairs of Electrons = 1
Bond Angle = 90°

31. Metallic Bonds How atoms are held together in the solid.
Metals hold onto there valence electrons very weakly.
Think of them as positive ions floating in a sea of electrons.

32. Sea of Electrons + Electrons are free to move through the solid.
Metals conduct electricity. +

33. Hammered into shape (bend). Ductile - drawn into wires.
Metals are Malleable
Hammered into shape (bend).
Ductile - drawn into wires.

34. Malleable +

35. Malleable
Electrons allow atoms to slide by. + + + + + + + + + + + +

36. EXIT QUIZ What bonds do we know? What is major difference?
Which the most unique?