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STOICHIOMETRY

UNIT 2. STOICHIOMETRY. Video 2.1. 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.

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STOICHIOMETRY

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  1. UNIT 2 STOICHIOMETRY

  2. Video 2.1 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. Video 2.2 Calculating Moles

  9. What if you have more than one mole of a substance? Use the formula on table T: number of moles = given mass (g) _ gram-formula mass (Given mass will be your answer.)

  10. Let’s try these together… • Calculate the mass of 6.70 moles of carbon. • Calculate the mass of 0.023moles of lithium. • Calculate the mass of 25.02 moles of calcium phosphate.

  11. The gfm formula • Use the same formula to calculate the moles, placing the number in the question on the numerator: number of moles = given mass (g) gram-formula mass

  12. Let’s try these together… 4. Determine the number of moles in 8.0 grams of Boron. • Determine the number of moles in 0.567 grams of Helium. • Find the number of moles in 1230 grams of magnesium sulfate.

  13. Video 2.3 Moles to Coefficients

  14. Coefficients 4Al + 3O2 2Al2O3 reactants products Coefficients: How many moles of the substance are needed in a reaction.

  15. Relating Moles • To relate moles of one substance to another, simply create a proportion: 4Al + 3O2 2Al2O3 • If 3 moles of oxygen react, how many moles of Aluminum oxide form? • If 4 moles of aluminum react, how many moles of oxygen are needed? • If 4 moles of aluminum oxide are formed, how many moles of oxygen were used? • If 8 moles of aluminum react, how many moles of oxygen are needed?

  16. Relating Moles 16 Al + 3S8 8Al2S3 • If 2.50 moles of sulfur react, how many moles of aluminum sulfide form? • If 3.75 moles of aluminum react, how many moles of sulfur are needed?

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

  18. Video 2.4 Balancing

  19. 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

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

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

  22. 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.

  23. Video 2.5 Empirical and Molecular Formulas

  24. Empirical Formulas • Determine the mass of each element in water. • What is the ratio of mass of hydrogen to mass of oxygen in water? • Is the ratio of mass related to the formula of water? • Find the moles of H and O in water. • Is the ratio of moles related to the formula of water?

  25. Empirical Formulas • Empirical formula refers to any molecular formula in it’s reduced form. Are these empirical? If not, reduce them: • C2H2 • C6H12O6 • NO2 • Na2(OH)2 CH CH2O NO2 NaOH

  26. Molecular Formulas • Molecular Formulas are some multiple of the empirical formula. • Example: 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?

  27. 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?

  28. Video 2.6 Percent Composition

  29. Percent Composition • Nutrition Facts on foods can tell you just how much of a substance you are consuming and how that relates to how much you should eat in a day. • It is equally important to know how much of a element or compound is in a mixture.

  30. Percent Composition • By mass: mass part x 100 total mass • Find the % by mass of phosphoric acid. • Find the percent by mass of Ca in Ca(OH)2.

  31. Find the percent by mass of Na: 23.0/74.5 *100 = 30.9% • NaClO • NaCl • NaOH 23.0/58.5 *100 = 39.3.% 23.0/40.0 *100 = 57.5%

  32. Percent Composition Examples 1. 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?

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