Try these..

1. Try these.. • Mg • CaCl2 • H2S

2. Double and Triple Bonds Example: HCN Make a table: atom have need H 1 2 C 4 8 N 58 total 10 18 Difference: 18-10=8 divide by 2 = 4 You need 4 bonds in this structure Sharing 4 or 6 electrons (Double or Triple bonds allow this to happen)

3. Electron dot drawings for polyatomic ions Always include brackets and charges, but have covalent bonds inside the ion Count the number of valence electrons for each and the add or subtract and electron to make the correct charge NH4+ OH- SO42- Draw NH4OH

4. Why is water so unique? Why Can water bugs run across a pond? Why does water have such a high boiling point? Why can we live on earth?

5. Why is water attracted to a – charge?Why is Hexane not attracted?

6. In this unit we will be able to understand how the Chemical Bonds in a substance determine physical properties • Why water is so unique • How the bonds that compose a substance determine the properties within • How shampoo works • How household cleaners work effectively

7. Lets set up your Lab Book • Purpose: To study the physical properties of common solids and to investigate the relationship between the type of bonding in a substance and its properties. • Volatility • Melting Point • Solubility • Brittleness • Conductivity

8. Procedure: See Handout • Volatility-Waft the substance • Solubility (Hexane and Water) in well plate • Conductivity (RED &GREEN LIGHT MEANS CONDUCTIVE) • Melting Point Watch Glass on beaker of water and test tube in bunsen burner • Brittleness (MORTAR AND PESTLE STATION)

9. Data • Record Observations in Table

10. Disposal • Rescue Aluminum if possible • Rinse out Sand in Garbage • Everything else can go down Sink

11. What did you discover in the periodic properties lab? • Which substance was the most volatile? • Which substances had the lowest melting point? • Which substances conducted electricity? • Which substances dissolved in water? Hexane? • Which substances do you believe had the strongest bonds? Why? • Which substances do you believe have the weakest bonds? Why?

12. How do you determine the types of bond that exist in a compound? • What is electronegativity? • For the following molecules,and ionic compounds • Draw the Lewis Dot Structure • CaS • AlCl3 • BH3

13. Electron dot drawings for polyatomic ions Always include brackets and charges, but have covalent bonds inside the ion Count the number of valence electrons for each and the add or subtract and electron to make the correct charge NH4+ OH- SO42- Draw NH4OH

14. Exceptions to the octet rule • Metals MgH2 BH3

15. 3. Some Nonmetal atoms because of their size, they can have more than an octet of electrons (due to the presence of empty “d” orbitals which can be used for bonding). SF6 PCl5 DON’T FOCUS ON THESE BUT KNOW THEY OCCUR!

16. Try these…. • Mg(OH)2 C3H6 O2

17. Note: • Not all covalent bonds have equal sharing of electrons • There are electron hogs!!! Elements that hold on to the electrons more tightly than others • You can determine if a bond is ionic,covalent and if there is an electron hogs, through looking at a characteristic property.

18. When the ΔE.N. is less than 2.0, the bond is covalent Examples: The O-H Bond in H2O The N-O Bond in NO2 • This means the electrons spend more time around one of the elements giving it a partial charge • Draw a picture of how you think the electrons would be distributed for an OH bond and a NO bond When the electrons are shared equally ex: H-H bond NCl Bond the bond is pure covalent and has no partial charge Draw a picture that describes what this would look like Why do you think there would not be a partial charge on these bonds?

19. The type of bond can usually be calculated by finding the difference in electronegativity of the two atoms that are going together.

20. Electronegativity Difference • If the difference in electronegativities is between: • 2.0 to 4.0: Ionic • Covalent Bonds • 0.1 to 1.9: Polar Covalent • 0.0: Non-Polar Covalent • Example: NaCl • Na = 0.8, Cl = 3.0 • Difference is 2.2, so • this is an ionic bond!

21. These bonds are called intramolecular forces • Have various strengths • Ionic (STRONGEST) • Polar Covalent (NEXT STRONGEST) • Covalent (STRENGHTH DEPENDS ON ELECTRONEGATIVITY DIFFERENCE)

22. Bond Polarity HCl is POLAR because it has a positive end and a negative end. (difference in electronegativity) Cl has a greater share in bonding electrons than does H. Cl has slight negative charge (-d) and H has slight positive charge (+ d)

23. Bond Polarity This is why you can dissolve Glucose in water and not hexane… Glucose is charged thus will be attracted to the charged bonds in water

24. Today we’re Putting it all together! • Why do substances have the properties that they do? • How can I predict the physical state of a substance?

25. Draw the Lewis Dot and Predict the shape for • I2 • SO42- • TeBr2

26. The difference between polar bonds and polar molecules • Polar bonds have a electronegativity difference between 2 atoms • Polar molecules • Include a polar bond • Have an asymmetrical shape • Is NH3 a Polar Molecule? • Yes! Why?

27. Is CCl4 a polar molecule? Why? Does it have a polar Bond? Yes Do it have an asymmetrical shape? NO!!! Thus it is not a polar molecule. What about CO2 Does it have a polar bond? Is it asymmetrical? NO Thus it is not a polar molecule..

28. Intermolecular Forces and Bonding in Solids

29. What holds molecules together?What holds several molecules together?How does this relate to Properties? Intramolecular Forces (Those bonds that hold molecules or compounds together internally) vs. Intermolecular Forces (Those forces that exist BETWEEN molecules).

30. Generally, intermolecular forces are much weaker than intramolecular forces. Intermolecular Forces Intermolecular forces are forces between molecules (Intermural sports are between different schools). Intramolecular forces hold atoms together in a molecule. Intramural sports are competition at a specific school. • Intermolecular vs Intramolecular • 41 kJ to vaporize 1 mole of water (inter) • 930 kJ to break all O-H bonds in 1 mole of water (intra) “Measure” of intermolecular force boiling point melting point

31. Types of intermolecular Forces • Dispersion Forces Weak Intermolecular Forces (In Non Polar Bonds) • Dipole-dipole interactions (In Polar Bonds) • Hydrogen bonds (Bonds between H and F,O,N Ex: H2O, NH3 • Relative strength of Intermolecular Forces: • Weakest to strongest: dispersion forces, dipole-dipole, hydrogen bonds • All are much weaker than intramolecular forces (covalent bonds,ionic bonds or metallic bonds)

32. What type of Intermolecular Forces exist • In Non Polar molecules • Dispersion Forces • In Polar Molecules • Dipole Dipole Forces, dispersion forces and sometimes Hydrogen Bonds

33. O O S Intermolecular Forces What type(s) of intermolecular forces exist between each of the following molecules? CH4 CH4 is nonpolar: dispersion forces. SO2 SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules. HF HF is a polar molecule: dipole-dipole forces. Hydrogen is bounded to F. Hydrogen bonds exist. There are also dispersion forces between HF molecules.

34. How does mass affect dispersion forces? Boiling Points of Group IV Hydrides

35. Which molecule will have a lower boiling point? AsH3 or PH3? Why?

36. Molecule Polarity • “Like Dissolves Like” • Polar dissolves Polar • Nonpolar dissolves Nonpolar

37. How do I put all this information together? • Determine if a Substance is a Ionic Compound or Molecule • If the substance is a molecule,determine if it is polar • Look the type of intermolecular forces in the molecule • Look at the Mass of the substance

38. Why does water have the qualities that it does? • What types of bonds? • What types of intermolecular forces at room temperature? • Why types of substances will dissove in water? • Will NaCl dissolve in water? • Will NaCl dissove in hexane? • Will Lauric acid dissolve in water? • Will Lauric acide dissolve in hexane?