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Stoichiometry

Stoichiometry. Chapter 12. What is stoichiometry? . The study of quantitative relationships between amounts used and products formed by a chemical reaction Based on the law of conservation of mass

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Stoichiometry

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  1. Stoichiometry Chapter 12

  2. What is stoichiometry? • The study of quantitative relationships between amounts used and products formed by a chemical reaction • Based on the law of conservation of mass • Chemical bonds in reactants break and new chemical bonds form to produce products, but the amount of matter present at the end of the reaction is the same as was present in the beginning • Mass of reactants = mass of products

  3. Mass of reactants = Mass of products Example: CH4 + 2O2 CO2 + 2H20 Mole Mass • CH4 • 2O2 • CO2 • 2H20

  4. Mole Ratios • Mole Ratio- a ratio between the number of moles of any substances in a BALANCED chemical equation Example: CH4 + 2O2 CO2 + 2H20 • 1 mol CH4 :2 mole O2 • 2 molH20 : 2 molO2 • Your Turn  Determine all possible mole ratios for the following chemical equation: • 3Fe(s) + 4 H20(l) Fe3O4(s) + 4 H2(g)

  5. 3Fe(s) + 4 H20(l) Fe3O4(s) + 4 H2(g)

  6. Practice Problems • Write a balanced chemical equation for each reaction and determine the possible mole ratios. • Nitrogen reacts with hydrogen to produce ammonia (NH3) • Hydrogen peroxide (H2O2) decomposes to produce water and oxygen. • Pieces of zinc react with a phosphoric acid solution to produce solid zinc phosphate and hydrogen gas.

  7. N2(g) + 3H2(g) 2 NH3(g) 2H2O2  2H2O + O2 3Zn(s) + 2 H3PO4(aq)  Zn3(PO4)2(aq) + 3 H2(g)

  8. Stoichiometric Calculations • Mole Mole Calculations • Plan to solve: • Balance Equation • Identifygiven quantity (in mol) • Conversion factor (mole ratio) • Calculate unknown quantity (in mol) Example: The elements lithium and oxygen react explosively to form lithium oxide, Li2O. How many moles of lithium oxide will form if 2 mol of lithium react?

  9. Mole Mass Calculations • Plan of solve: • Balance Equation • Identifygivenquantity (in mol) • Conversion factor (mole ratio) • Molar mass of unknownquantity (in g/mol) • Calculate unknownquantity (in grams) • What mass, in grams, of glucose is produced when 3.00 mol of water reacts with carbon dioxide? __CO2 (g) + __H2O(l)__C6H12O6 (s) + __O2(g)

  10. Mass Mole Calculations • Plan of attack: • Balance Equation • Identifygivenquantity (in grams) • Molar mass of givenquantity (in g/mol) • Conversion factor (mole ratio) • Calculateunknownquantity (in mol) • The first step in the industrial manufacture of nitric acid is the catalytic oxidation of ammonia. The reaction is run using 824g NH3 and excess oxygen. How many moles of NO are formed? __NH3(g) + __O2 (g) __NO(g) + __H2O (g)

  11. Mass Mass Calculations • Plan of attack: • Balance Equation • Identifygiven quantity (in grams) • Molar mass of givenquantity (in g/mol) • Conversion factor (mole ratio) • Molar mass of unknown quantity (in g/mol) • Calculate unknownquantity (in grams) • How many grams of SnF2 are produced from the reaction of 30.00g of HF with Sn?

  12. Practice Problem  • Sulfuric acid is formed when sulfur dioxide reacts with oxygen and water. Write the balanced chemical equation for the reaction. • If 12.5 mol SO2 reacts, how many moles of H2SO4 can be produced? • If 2.50 g SO2 react with excess oxygen and water, how many grams of H2SO4 are produced?

  13. 2SO2(g) + O2(g)+ 2H2O(l)  2H2SO4(aq)

  14. Limiting Reactants • Why do reactions stop? • How many complete cars can you make?

  15. Limiting Reactant - The reactant in a chemical reaction that limits the amount of product that can be formed.  The reaction will stop when all of the limiting reactant is consumed. • Excess Reactant - The reactant in a chemical reaction that remains when a reaction stops when the limiting reactant is completely consumed.  The excess reactant remains because there is nothing with which it can react.

  16. The “Have and Need” Method • How can you determine which reactant is limited? • Masses of both reactants are given • Determine the grams or moles of both reactants given (“Have”) • Determine the grams or moles “needed” of each reactant by relating them

  17. Practice Problem • When 36.0 g of H2O is mixed with 167g of Fe, which is the limiting reactant? • What mass in grams of black iron oxide is produced? • What mass in grams of excess reactant remains when the reaction is completed? • 3Fe (s) + 4H2O(g) Fe3O4(s) + 4H2 (g) • Determine the grams of both reactants given (“Have”) • Determine the grams “needed” of each reactant by relating them with their “Have and Need”

  18. 3Fe (s) + 4H2O(g) Fe3O4(s) + 4H2 (g) • 36.0 g H20 • 167 g Fe • Use the limiting reactant to calculate the product mass

  19. Addition Practice 6 Na(s) + Fe2O3(s) 3Na2O (s) + 2 Fe (s) If 100.0g Na and 100.0g Fe2O3are used in the reaction, determine • The limiting reactant • The reactant in excess • The mass of solid iron produced • The mass of excess reactant that remains after the reaction is complete

  20. 6 Na(s) + Fe2O3(s) 3Na2(s) + 2 Fe (s)100.0g Na and 100.0g Fe2O3

  21. Percent Yield • The theoretical yield is the maximum amount of product from a given amount of reactant • The actual yield is the amount of product actually produced when the chemical reaction is carried out in an experiment • The percent yield of product is the ratio of the actual yield to the theoretical yield expressed as a percent % Yield = actual yield (from an experiment) x 100 theoretical yield (from calculation)

  22. Example • If 75.0g of CO reacts to produce 68.4g of methanol (CH3OH), what is the percent yield of CH3OH? • CO(g) + 2H2 (g) CH3OH (l)

  23. 2Al(s) + 3CuSO4(aq)3Cu (s) + Al2 (SO4)3 (aq) • Aluminum reacts with an aqueous solution containing excess copper (II) sulfate. If 1.85 g Al reacts and the percentage yield of Cu is 56.6%, what mass of Cu is produced?

  24. Use the data to determine the percent yield of the following reaction. • 2Mg (s) + O2(g) 2MgO (s)(Oxygen is in excess)

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