1. Chemistry Review: day 2 Get a packet and a calculator

2. Stuff: Reviews: Reviews: –Today: Unit 3 and Unit 4 –Friday: Unit 5 and Unit 6 –Powerpoints on website –Next Tuesday afternoon – final weak points & equilibrium –Tutoring: Monday and Tuesday afternoon. Pizza Tuesday. –Wednesday, June 8 th – final.

3. Today: Review yesterday’s stuff Review yesterday’s stuff Bonds, Naming, Equations, Stoichiometry Bonds, Naming, Equations, Stoichiometry Lunch: 11:00 – 11:20 Lunch: 11:00 – 11:20 Quiz 2: Right after lunch (don’t be late!) Quiz 2: Right after lunch (don’t be late!) Movie continues: Big Hero 6 Movie continues: Big Hero 6

4. Chemical Bonds  Transferring e -  Sharing e - IONIC (metal + nonmetal) COVALENT (nonmetal + nonmetal)

5. Formulas: Ionic Criss-Cross: the absolute value of the oxidation number of an element becomes the opposite element’s subscript. Criss-Cross: the absolute value of the oxidation number of an element becomes the opposite element’s subscript. Metal comes first, followed by the nonmetal. Metal comes first, followed by the nonmetal.

7. Naming: Metal/cation’s name first, then the nonmetal/anion Metal/cation’s name first, then the nonmetal/anion Replace the nonmetals end letters with –ide (- ine, -gen, -ur, -orous change to -ide) Replace the nonmetals end letters with –ide (- ine, -gen, -ur, -orous change to -ide) Never ever change the name of a polyatomic ion, whether it’s first or last. Never, ever. Never ever change the name of a polyatomic ion, whether it’s first or last. Never, ever.

8. Covalent Naming: Prefixes in name tell you the formula. Prefixes in name tell you the formula. Example: Carbon dioxide = CO 2 (one carbon and two oxygen)

10. Polar Molecules: Only happens in covalently bonded molecules. “Polar” means that the molecule has “poles” or a (+) area and a (-) area. “Polar” means that the molecule has “poles” or a (+) area and a (-) area. A molecule’s polarity is due to the unequal distribution of electrons as they are pulled toward the more electronegative element. A molecule’s polarity is due to the unequal distribution of electrons as they are pulled toward the more electronegative element.

11. Water is a polar molecule Can “rip” ionic compounds like NaCl apart. Can “rip” ionic compounds like NaCl apart. Positive part of H 2 O pulls at negative Cl -, negative part of H 2 O pulls at positive Na +. Positive part of H 2 O pulls at negative Cl -, negative part of H 2 O pulls at positive Na +.

12. I say: You say Ionic: salt Ionic: salt Covalent: sugar Covalent: sugar Polar: water Polar: water

13. The longer the covalent bond, the less strength. BONDLENGTHSTRENGTH Single1 (longest)3 Double22 Triple31 (strongest)

14. VSEPR Theory Valence Shell Electron Pair Repulsion Electron pairs want to be as far away as possible.

15. Linear Ex: Ex: HClHCl HClHCl One bond between 2 elements, regardless of bond and number of lone pair e- forms a shape called LINEAR.

16. Tetrahedral Tetrahedral H Ex: Ex: CH 4 H C H CH 4 H C H H Central atom (C) contains four bonds and no lone pair of e-. Everyone pushes everyone else equally. Central atom (C) contains four bonds and no lone pair of e-. Everyone pushes everyone else equally.

17. Pyramidal Ex: Ex: NH 3 HNH NH 3 HNHH The central atom (N) has three bonds and one lone pair of e-. The lone pair pushes “harder” than the shared pairs. The central atom (N) has three bonds and one lone pair of e-. The lone pair pushes “harder” than the shared pairs.

18. Bent Triatomic Ex: Ex: H 2 O H 2 O HO HOH The central atom (O) has two bonds and two lone pair of e-. They push the shared pairs. The angle between the shared pairs in water is 105 o. The central atom (O) has two bonds and two lone pair of e-. They push the shared pairs. The angle between the shared pairs in water is 105 o.

19. Linear Triatomic Ex: Ex: CO 2 OCO CO 2 OCO The central atom (C) has two double bonds, or one triple bond and a single bond, and no lone pair e-. Double shared bonds DO NOT MOVE. THEY’RE STUCK. They keep the shape. Bond angle is 180 o. The central atom (C) has two double bonds, or one triple bond and a single bond, and no lone pair e-. Double shared bonds DO NOT MOVE. THEY’RE STUCK. They keep the shape. Bond angle is 180 o.

20. Trigonal Planar (showing resonance) Ex: Ex: SO 3 SO 3 OSO O The central atom (S) has 4 bonds, and no lone pair e-. One of the oxygens has a double bond. The central atom (S) has 4 bonds, and no lone pair e-. One of the oxygens has a double bond.

21. Would this be Polar?

22. No. Think about where the electrons are…

23. What about Ammonia? Yes. The electrons are not symmetrically distributed around the molecule. (Remember the lone pair up top?)

24. How do I know? Electronegativity Difference Range Most Probable Type of Bond Example 0.0 – 0.4Nonpolar covalentH-H (diatomics) 0.4 – 1.0Moderately polar covalent H-Cl (0.9) 1.0 – 2.0Very polar covalentH-F (1.9) > 2.0IonicNaCl

25. Relative strength of bonds, chemical bonds Intramolecular Forces (within the molecule – ionic/covalent) Ionic>Triple Covalent> Double Covalent> Single Covalent

26. InTRA vs InTER

27. Intramolecular forces keep the molecule together. Intramolecular forces keep the molecule together. Intermolecular forces hold many molecules together. They help to determine if you’re a solid, liquid or a gas. If you don’t have a strong connection to your buddies – you’ll move away. Intermolecular forces hold many molecules together. They help to determine if you’re a solid, liquid or a gas. If you don’t have a strong connection to your buddies – you’ll move away.

28. 1) London Dispersion Force (weakest) Happens between “momentary dipoles” in molecules, when electrons are momentarily more concentrated in one area over another. Seen in non-polar molecules. Also in CCl 4. Happens between “momentary dipoles” in molecules, when electrons are momentarily more concentrated in one area over another. Seen in non-polar molecules. Also in CCl 4.

29. 2) Dipole-Dipole – between polar molecules. + area of one is attracted to – area of the other.

30. 3) Hydrogen Bonds (strongest) Happen between covalently bonded hydrogen and a highly electronegative atom Happen between covalently bonded hydrogen and a highly electronegative atom

31. The type of intermolecular bonds you have…  Tell me what happens if you have strong hydrogen bonds. What are you at room temperature? Liquid or solid (higher boiling point)  What about dipole-dipole?  What about dipole-dipole? Liquid -> gas  What about if your molecules are bonded by dispersion forces?  What about if your molecules are bonded by dispersion forces? Gas

32. Macromolecules Large molecule groups formed by covalent or intermolecular forces between smaller, similar molecules. Large molecule groups formed by covalent or intermolecular forces between smaller, similar molecules. Structure of molecules and their IMFs lead to bigger and bigger molecules and collections of molecules. Structure of molecules and their IMFs lead to bigger and bigger molecules and collections of molecules.

33. Network solids All of the atoms are covalently bonded to one another in a 2D or 3D network. Examples: diamonds, quartz All of the atoms are covalently bonded to one another in a 2D or 3D network. Examples: diamonds, quartz

34. Polymer – covalently bonded repeating small units PVC – dipole-dipole between C and Cl PVC – dipole-dipole between C and Cl

35. The Mole A mole (mol) of anything is Avogadro’s number: 6.022 x 10 23 of that stuff. A mole (mol) of anything is Avogadro’s number: 6.022 x 10 23 of that stuff. A mole of a substance weighs the atomic mass of that substance, expressed in grams A mole of a substance weighs the atomic mass of that substance, expressed in grams For a compound: a mole is the formula (or molecular) mass in grams/mole. For a compound: a mole is the formula (or molecular) mass in grams/mole.

36. Formula mass is the same as molar mass or gram formula mass, but with different units. Instead of amu, the unit is g/mole Ex: NaCl  58.44 g/mole Ex: H 2 O  18.01 g/mole

37. And… Moles -> mass Moles -> mass Mass -> moles Mass -> moles Mass -> moles -> particles Mass -> moles -> particles Etc… Etc…

38. Example Example 8.76 mole of H 2 O = ? molecules H 2 O 8.76 mole of H 2 O = ? molecules H 2 O 8.76 moles H 2 O 6.02 x 10 23 molecules = 8.76 moles H 2 O 6.02 x 10 23 molecules = 1 1 mole 1 1 mole 5.27 x 10 24 molecules 5.27 x 10 24 molecules

39. Molar Volume  the volume occupied by one mole of a gas at STP is 22.414 L Molar Volume  the volume occupied by one mole of a gas at STP is 22.414 L 22.414 L/mole  MUST BE AT STP (standard temperature and pressure) 22.414 L/mole  MUST BE AT STP (standard temperature and pressure)

41. Percent Composition The % (by mass) of each element in a compound: The % (by mass) of each element in a compound: (mass of element in the compound/mass of the compound) x 100 % = percent mass of compound For example, H 2 O… For example, H 2 O…

42. Water Example: H 2 O  GFM = 18.02g / mole H  GFM = 1.0 g/mole % composition H = 2(1.0 g)/18.02 g x 100% Answer  Water is 11.10 % Hydrogen % composition O = (16.0 g/ 18.02 g) x 100 % Answer  Water is 88.89 % Oxygen

43. Grams… How many grams of oxygen are in 25g of H 2 O? How many grams of oxygen are in 25g of H 2 O? What % of H 2 O is oxygen? (from the last slide…) – % Oxygen in H 2 O = 88.89 % –.89 (25g) = 22.25g of oxygen

44. Empirical & Molecular Formulas Empirical Formula – the formula for a compound expressed as the smallest possible whole-number ratio of subscripts of the elements in the formula. For example, C 2 H 3 Molecular Formula- the formula for a compound in which the subscripts give the actual number of each element in the formulas as it truly exists. For example, C 6 H 9

45. First - Calculating Empirical Formulas Remember: Remember: Percent to mass Percent to mass Mass to mole Mass to mole Divide by small Divide by small Multiply ‘til whole Multiply ‘til whole

46. Calculating Molecular Formulas Given molecular weight/empirical GFM = molecular formula multiplier Given molecular weight/empirical GFM = molecular formula multiplier The molecular formula can be identical to the empirical formula, but doesn’t have to be. The molecular formula can be identical to the empirical formula, but doesn’t have to be.

47. 2H 2 + O 2 -> 2H 2 O Why does this work?

48. How many hydrogens and oxygens on the left and right? 2H 2 + O 2 -> 2H 2 O

49. 4 hydrogens 2 oxygens 4 hydrogens 2 oxygens

50. In (more) other words…

51. 5 types of chemical reactions to know: 5 types of chemical reactions to know: 1.Synthesis (A + B  AB) 2.Decomposition (AB  A + B) 3.Single Replacement (AB + C  AC + B) 4.Double Replacement (AB + CD  AD + CB) 5.Combustion (C n H n + O 2  CO 2 + H 2 O) If you know what type of reaction you have, you can predict the product(s).

52. Moles from equations… “how many grams of ____ are produced when…” Balance equation Balance equation Start with given, change to moles Start with given, change to moles Molar ratio (from balanced equation) Molar ratio (from balanced equation) Change back to what you need (if not moles) Change back to what you need (if not moles)