1 / 65

Unit 1

Moles and Dimensional Analysis. Unit 1. Video AP1.1. Moles and Molar Mass. Stoichiometry. The Mole represents a specific amount of any substance. Specifically it represents 6.02x10 23 particles such as atoms and molecules. The Mole = 6.02x10 23.

lionel-franklin
Download Presentation

Unit 1

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Moles and Dimensional Analysis Unit 1

  2. Video AP1.1 Moles and Molar Mass

  3. Stoichiometry • The Mole represents a specific amount of any substance. Specifically it represents 6.02x1023 particles such as atoms and molecules.

  4. The Mole = 6.02x1023 • The mole is based on the fact that 12 grams of Carbon-12 has a mole of atoms. • A mole means you have 6.02x1023 particles.

  5. Molar Mass • Molar Mass (aka gram formula mass or molecular mass) is the mass of one mole of a substance. • Element’s molar masses are reported on the periodic table.

  6. Molar Mass Examples: Elements • What is the molar mass of iron? • What is the molar mass of copper? • 55.8 g/mol • 63.5 g/mol

  7. Molar Mass Examples: Compounds • What is the molar mass of water? • H2O = 2(1.0) + 16.0 = • 18.0g/mol • What is the gram-formula-mass of calcium chloride? • CaCl2 = 40.1 + 2(35.5) = • 111.1 g/mol

  8. Dimentional Analysis • Calculate the mass of 6.70 moles of iodine.

  9. Dimensional Analysis 2. Calculate the mass of 0.023moles of NH3.

  10. Dimensional Analysis 3. Determine the number of moles in 8.0 grams of CCl4.

  11. Dimensional Analysis 4. Determine the number of moles in 0.567grams of Helium.

  12. DA pointers summarized • Start with the number given as a numerator over 1. • MULTIPLY by a new fraction • Place a conversion with the same unit on the denominator of the new fraction. • Multiply numerators, divide denominators.

  13. Video AP 1.2 Molar Volume and Avogadro

  14. Mole Road Map

  15. Examples 1. Calculate the number of molecules in 4.37 grams of Lithium sulfide. 2. Calculate the number of atoms in question 1.

  16. Examples 3. A gas at standard temperature and pressure will have a molar volume of 22.4L. What is the volume of 7.89moles gas at STP?

  17. Examples 4. How many atoms are present in 79.80 Liters of diphosphorouspentoxide gas at STP?

  18. THINK Why do chemists use moles to measure substances? Why aren’t grams, liters and molecules enough?

  19. Video AP 1.4 Theoretical and Percent Yield

  20. Limiting Reagents • Given: 20 11 36 • 1 sandwich = 2 + 1 +3 • How many sandwiches can I make if I used all of the bread with unlimited other resources? • How many sandwiches can I make if I used all of the cheese with unlimited other resources?

  21. Limiting Reagents • Given: 20 11 36 • 1 sandwich = 2 + 1 +3 3. How many sandwiches can I make if I used all of the ham with unlimited other resources? 4. How many sandwiches can I make if I used only what is given? 5. What ingredient was the limiter? Why?

  22. Limiting Reagents Examples 1. S + 3F2 SF6 Suppose you have 4 moles of sulfur and 2 moles of fluorine, which is the limiter and which is in excess? How many moles of SF6 can be produced? 4S produces 4 SF6 2F2produces 0.6 moles of SF6 (F2 Limiter!)

  23. Limiting Reagents Examples 2. 2NH3 + CO2 (NH2)2CO + H2O If 637.2 grams of ammonia react with 1142 grams of carbon dioxide, find the limiter and the amount of urea produced. NH3produces 1123 grams of urea (NH3 Limiter!) CO2 produces 1559 grams of urea.

  24. Limiting Reagents Examples 3. In question 2, how much of the excess is left over after the reaction goes to completion? If 1123 grams are produced, go backwards to find that 823 grams of CO2 were used: 1142-823 = 319 g CO2 left over

  25. Limiting Reagents Examples 4. 5Ca + V2O5 5CaO + 2V 1.96x103 grams of calcium react with 1.54x 103 grams of V2O5. Find the limiter and the amount of V produced. 5. If 803 grams of V are actually made calculate the percent yield.

  26. Video AP 1.5 Percent Compositon, Empirical and Molecular Formulas

  27. Percent Composition % By mass: mass part x 100 total mass • Find the % by mass of phosphoric acid. H3PO4 3(1.0) x 100 3(1.o)+32.1+4(16.0) = 3.02%

  28. Example 2 Find the percent by mass of Ca in Ca(OH)2. 40.0 x 100 40.0+2(16.0)+2(1.0) = 54.1%

  29. Percent Composition Examples A sample of a substance containing only magnesium and chlorine was tested in the laboratory and found to be composed of 74.5% chlorine by mass. If the total mass of the sample was 190.2 grams, what is the mass of the magnesium? 0.745 (190.2) = • 141.7 grams

  30. Empirical Formulas • Empirical formula refers to any molecular formula in it’s reduced form. Are these empirical? If not, reduce them: • C4H8 • C6H12O6 • N3O6 • Na2(OH)2 CH2 CH2O NO2 NaOH

  31. Molecular Formulas • Molecular Formulas are some multiple of the empirical formula. • If the empirical formula is CH4 a molecular formula could be CH4, C2H8, C3H12 etc. • A compound whose empirical formula is NH3 has a mass of 34 g/mol. What is the empirical formula? NH3 = 17.0 g/mol 34/17 = 2 2NH3 = N2H6

  32. Steps: • Find the mass of the empirical formula. • Divide the mass given by the empirical mass. • Distribute your answer through the empirical formula. • If a compound has a mass of 45 g/mol and an empirical formula of CH3, what is the molecular formula? CH3 = 15 g/mol 45/15 = 3 3CH3 = C3H9

  33. Calculating Empirical Formulas From Percent Composition Rules: Find the empirical formula of a compound with 40.92% C, 4.58% H and 54.50% O by mass. • Convert the % to grams, assuming the total mass of the sample is 100 grams. • 40.92g C, 4.58g H, 54.50g O • 2. Convert mass to moles, using atomic masses. • 40.92/12 = 3.40 moles C • 4.58/ 1 = 4.58 moles H • 54.50/16 = 3.4o moles O • Divide each element’s moles by the smallest number of moles. • Find a whole number ratio. • 1 mole C, 1.5 mole H, 1 mole O • C2H3O2

  34. Now try these… • A sample of nitrogen and oxygen contains 1.52 g N and 3.47g O and a molar mass of about 90 grams. Find the empirical and molecular formulas.

  35. Example 2 Find the empirical and molecular formula of a 30 gram sample with 6.44 grams of Boron and 1.803 grams of hydrogen.

  36. Multiple Choice A sample of a compound contains 47.0% silicon and the rest is oxygen. If the total mass is 180 grams, what is a possible set of empirical molecular formulas of the compound? a. SiO2 and SiO b. Si4O4 and SiO c. Si2O4 and SiO2 d. Si3O6 and SiO2

  37. Video AP 1.6 Balancing and Types of Reactions

  38. Conservation of Mass • In a reaction, atoms and molecules cannot appear or disappear. Mass must stay constant from the beginning to the end of the reaction. • H2 + O2  H2O • ___H2 + ___O2  ___ H2O

  39. Balancing Reactions ___ N2 + ____H2  ____ NH3 ___Li + ____O2  ___Li2O

  40. Balancing __Pb(NO3)2 +__K2CrO4___PbCr2O4 + ___KNO3 ___C4H8 + ___O2  ___CO2 + ___H2O

  41. Types of Reactions • Synthesis: A + 2B  AB2 • Decomposition: AB2 A + 2B • Combustion: CH4 + O2 CO2 + H2O • Single Replacement: AB + C  CB +A • Double Replacement: AB + CD  AD + CB * Notice synthesis and decomposition are opposites. Also, combustion can have any carbon compound as a reactant.

  42. Examples: • Determine the type of reaction: • NaCl + LiBr LiCl + NaBr • Mg + H2O  H2 + Mg(OH)2 • C3H8 + 5O2  3CO2 + 4H2O • S + 3F2  SF6 • N2H4 N2 + 2H2 DR SR C S D

  43. Video AP 1.7 Molarity and Dilutions

  44. Concentration • Concentrated solutions contain large amounts of solutes dissolved in the solvent. • Dilute solutions contain small amounts of solutes dissolved in solvent. Which is concentrated?

  45. Molarity • The amount of solute (dissolvable substance) in a solvent (does the dissolving like water) is known as concentration and can be represented many ways. • Molarity is the moles of solute per liter of solution. Find this on your reference tables! • The units for molarity are mol/L ormol.L-1or just M.

  46. mol of solute L of solution M = Molarity (M) • Because volume is temperature dependent, molarity can change with temperature.

  47. Examples can be done without a calculator! • If 3 moles of NaCl are dissolved in 6L of water, what is the molarity? • If 29 grams of NaCl are dissolved in one liter of water, what is the molarity? 3moles/6L = 0.5M 29g/58g = 0.5moles 0.5moles/1L = 0.5M

  48. Examples may need a calculator or can be estimated • If 100 grams of KF are dissolved in 300ml of water, what is the molarity? • Calculate the volume needed to create a 2M solution with 3.5 moles of Li2O. 100g/58.1g =1.72mol 1.72mol/.3L = 5.7M X = 1.75L 2M=3.5mol/x

  49. Try these… (without a calc!) 5. A solution is prepared with 4.o moles of NaCl in 2.oL water. Calculate the molarity. 6. A solution contains 202g of KNO3 (gfm=101) in 4.oL of solution. Calculate the molarity. 7. A 37g sample of KCl (gfm=74) is dissolved in 250mL of water. Calculate the molarity. 2.0M 0.5M 2.0M

  50. THINK If you have 3.12g of HCl and 250mL of H2O, what is the molarity of the acid? If 250mL of H2O is added again, how does the molarity change? M decreases by half

More Related