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Wednesday, Oct. 2 nd : “A” Day Thursday, Oct. 3 rd : “B” Day Agenda

Wednesday, Oct. 2 nd : “A” Day Thursday, Oct. 3 rd : “B” Day Agenda. Sec 9.2: “Limiting Reactants and Percentage Yield” Limiting/ Excess Reactants, Theoretical Yield, Actual Yield, Percentage Yield Homework: Sec. 9.2 review, pg. 319: #1-10 Concept Review.

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Wednesday, Oct. 2 nd : “A” Day Thursday, Oct. 3 rd : “B” Day Agenda

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  1. Wednesday, Oct. 2nd: “A” DayThursday, Oct. 3rd: “B” DayAgenda • Sec 9.2: “Limiting Reactants and Percentage Yield” • Limiting/ Excess Reactants, Theoretical Yield, Actual Yield, Percentage Yield • Homework: • Sec. 9.2 review, pg. 319: #1-10 • Concept Review

  2. Limiting Reactants and Theoretical Yield • In the previous section, we assumed that 100% of the reactants changed into products. • Theoretically, that is what SHOULD happen, but in the real world, other factors like these can limit the yield of a reaction: • Amounts of all reactants • Completeness of the reaction • Products lost in the process

  3. S’More Activity • See how many s’mores you can make from the marshmallows, graham crackers and chocolate that you’re given. • Which ingredient did you run out of first? • Which ingredients do you have left over? • In this case, the chocolate is the limiting reactant and the marshmallows and graham crackers are the excess reactants.

  4. The Limiting Reactant Forms the Least Amount of Product • Limiting Reactant: the substance that controls the quantity of product that can form in a chemical reaction. • This reactant is completely used up in the reaction. • Excess Reactant: the substance that is not used up completely in the reaction. • There will be some of this reactant left over after the reaction.

  5. The Limiting Reactants are often the More Expensive Reactants • In industry, the cheaper reactants are often used as the excess reactants. • In this way, the more expensive reactants are completely used up, saving the company money.

  6. Example One way to make hydrogen gas, H2 is: Zn + 2 HCl ZnCl2 + H2 • Question: If you combine 0.23 mol Zn and 0.60 mol HCl, would they react completely? • 0.23 mol Zn ? Mol H2 0.23 mol H2 • 0.60 mol HCl ? Mol H2 0.30 mol H2

  7. Example Zn + 2 HCl ZnCl2 + H2 0.23 mol Zn 0.23 mol H2 0.60 mol HCl 0.30 mol H2 • Since 0.23 mol of Zn makes less H2 than 0.60 mol of HCl, Zn is the limiting reactant and will be completely used up. • HCl is the excess reactant, meaning that there will be some HCl left over.

  8. Determine Theoretical Yield From Limiting Reactant • Theoretical Yield: the maximum quantity of product that a reaction could theoretically make if everything about the reaction works perfectly. • The theoretical yield is ALWAYS based on the limiting reactant.

  9. Sample Problem E, pg. 314 • Identify the limiting reactant and the theoretical yield of phosphorous acid, H3PO3, if 225 g of PCl3 is mixed with 123 g of H2O. PCl3 + 3 H2O H3PO3 + 3 HCl • Use stoichiometry to calculate the MASS of H3PO3 you could form from each reactant. • The reactant that produces the least amount of H3PO3is the limiting reactant. • The theoretical yield is the amount of H3PO3 produced from the limiting reactant.

  10. Practice PCl3 + 3 H2O H3PO3 + 3 HCl • Identify the limiting reactant and theoretical yield (in grams of HCl) for these reactants: 3.00 mol PCl3 and 3.00 mol H2O 3.00 mol PCl3 X 3 mol HCl X 36.46 g HCl = 328 g HCl 1 mol PCl3 1 mol HCl 3.00 mol H2O X 3 mol HClX 36.46 g HCl= 109 g HCl 3 mol H2O 1 mol HCl H2O is the limiting reactant. Theoretical yield is 109 g HCl.

  11. Actual Yield • Actual yield: The measured amount of a product of a reaction. • What you actually got from the experiment. • Theoretical yield = what you should get if everything works perfectly • Actual yield = what you actually get

  12. Actual Yield The actual yield is usually less than the theoretical yield. WHY? • Many reactions do not completely use up the limiting reactant. Some of the products turn back into reactants (reversible reaction). • The final product may need to go through additional purification processes (filtering, distilling, etc) and some product may be lost. • There could be other side reactions that use up the limiting reactant without making the desired product.

  13. Percentage Yield • The percentage yield is found by simply dividing the actual yield by the theoretical yield and multiplying by 100%. Percentage Yield = Actual Yield X 100% Theoretical Yield • The percentage yield describes the efficiency of a reaction.

  14. Sample Problem F, pg. 317 Determine the limiting reactant, the theoretical yield, and the percentage yield if 14.0 g of N2 are mixed with 9.0 g H2 and 16.1 g NH3 form. N2 + 3 H2 2 NH3 • Use stoichiometry to calculate the MASS of NH3 you could form from each reactant. • The reactant that produces the least amount of NH3is the limiting reactant. • The theoretical yield is the amount of NH3produced from the limiting reactant. • Use the theoretical yield and the actual yield (16.1 g) to calculate the percentage yield.

  15. Practice Determine the limiting reactant and the percentage yield for the following: N2 + 3 H2 2 NH3 14.0 g N2 react with 3.15 g H2 to give an actual yield of 14.5 g NH3. N2 is the limiting reactant. Theoretical yield = 17.0 g NH3. Percentage yield = 85.3%

  16. Homework • Sec. 9.2 review, pg. 319: #1-10 • Concept Review Quiz over section 9.2 next time…

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