BELLRINGER

1. BELLRINGER Explain in complete sentences the difference between polar molecules

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 Polaritysimulation 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 • 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. d+d- d+d- H F H F 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.

14. d+d- d+d- d+d- d+d- d+d- d+d- d+d- Dipole Interactions 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. d- d+ O d+ H d+ d- H H O H d+ Hydrogen Bonding

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

18. MOLECULAR SHAPES OF COVALENT COMPOUNDS

19. VSepR tHEORY ALENCE HELL VSEPR LECTRON AIR EPULSION

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

24. Trigonal Planar 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 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 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 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 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. Trigonal bIPyramidal 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 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. 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?