CHAPTER 9 OUTLINE STOICHIOMETRY
		I. The Arithmetic of Equations      A. Using Everyday Equations           1. Stoichiometry -- calculation of quantities in chemical reactions                a. Like bookkeeping or a recipe                b. How much product do you get from this amount of reactant?      B. Interpreting Chemical Equations                N2(g)  +  3H2(g)  -->  2NH3(g)           1. Particles                a. One molecule of nitrogen reacts with 3 molecules of hydrogen to form 2 molecules of ammonia                b. 1:3:2 ratio                c. Avogadro's # of nitrogen react with 3X Avogadro's # of hydrogen to form 2X Avogadro's # of ammonia           2. Moles                a. One mole of nitrogen reacts with three moles of hydrogen to form 2 moles of ammonia                b. Coefficients indicate # of moles           3. Mass                a. Must follow law on conservation of mass                b. 1 mol N2 = 28 g N2, 3 mol H2 = 6 g H2, 2 mol NH3 = 34 g NH3                c. 28 g N2  +  6 g H2  =  34 g NH3           4. Volume                a. Assume it is at STP (Standard Temperature and Pressure)                b. 22.4 L of nitrogen reacts with 67.2 L (3 X 22.4 L) of hydrogen to form 44.8 L (2 X 22.4 L) of ammonia II. Chemical Calculations      A. Mole-Mole Calculations           Example:      B. Mass-Mass Calculations           Example:      C. Particle-Particle Calculations           Example:      D. Volume-Volume Calculations           Example:      E. Mixed Calculations           Example: III. Limiting Reagent and Percent Yield      A. What is a Limiting Reagent?           1. Limiting reagent -- limits or determines the amount of product that can be formed           Example:      B. Calculating the Percent Yield           1. Usually in the lab, the product produced is less than what was expected based on the balanced chemical equation           2. Theoretical yield -- maximum amount of product that could be formed                a. Calculated from a balanced chemical equation           3. Actual (Experimental) yield -- amount of product that actually forms                a. Calculated from the lab experiment           4. Percent yield -- ratio of actual yield to the theoretical yield (in %)                     % yield  =  (actual yield/theoretical yield)  X  100%           5. Always less than 100%, Why?                a. Reaction does not go to completion                b. Impure reactants                c. Competing side reactions (unwanted products)                d. Loss of product during transfer                e. Not carefully measured Outline based upon:      Matta, M. S., Staley, D. D., Waterman, E. L., & Wilbraham, A. C. (2000). Chemistry, Addison-Wesley. (5th ed.). Menlo Park, CA: Prentice Hall, pp. 237-259.                                                              Return Home