Transcript Section 9.2
Wednesday, Oct. 2nd: “A” Day Thursday, Oct. 3rd: “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 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 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. 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. 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. 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 Example Zn + 2 HCl 0.23 mol Zn 0.60 mol HCl ZnCl2 + H2 0.23 mol H2 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. 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. 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 1. Use stoichiometry to calculate the MASS of H3PO3 you could form from each reactant. 2. The reactant that produces the least amount of H3PO3 is the limiting reactant. 3. The theoretical yield is the amount of H3PO3 produced from the limiting reactant. 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 HCl X 36.46 g HCl = 109 g HCl 3 mol H2O 1 mol HCl H2O is the limiting reactant. Theoretical yield is 109 g HCl. 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 Actual Yield The actual yield is usually less than the theoretical yield. WHY? 1. Many reactions do not completely use up the limiting reactant. Some of the products turn back into reactants (reversible reaction). 2. The final product may need to go through additional purification processes (filtering, distilling, etc) and some product may be lost. 3. There could be other side reactions that use up the limiting reactant without making the desired product. 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. 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 1. Use stoichiometry to calculate the MASS of NH3 you could form from each reactant. 2. The reactant that produces the least amount of NH3 is the limiting reactant. 3. The theoretical yield is the amount of NH3 produced from the limiting reactant. 4. Use the theoretical yield and the actual yield (16.1 g) to calculate the percentage yield. 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% Homework Sec. 9.2 review, pg. 319: #1-10 Concept Review Quiz over section 9.2 next time…