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Section 9.1 Using Chemical Equations Section 9.1 Using Chemical Equations 1984, Union Carbide Plant – 10000+ killed in Bhopal, India by Methyl Isocyanate 1988 Pepcon, Henderson Rocket Fuel Explosion Section 9.1 Using Chemical Equations Objectives 1. To understand more of the information given in a balanced equation 2. To use a balanced equation to determine relationships between moles of reactant and products Section 9.1 Using Chemical Equations A. Information Given by Chemical Equations • A balanced chemical equation gives relative numbers of reactant and product molecules (or moles) that participate in a chemical reaction. • The coefficients of a balanced equation give the relative numbers of molecules (or moles). CO(g) + 2H2(g) CH3OH(l) How many moles of CH3OH will be produced from 1, 2, or 5 moles of CO? How many moles of hydrogen will react with 1, 2, or 5 moles of CO? Section 9.1 Using Chemical Equations B. Mole-Mole Relationships in a BALANCED Reaction • A balanced equation can predict the moles of product that a given number of moles of reactants will yield. What number of moles of O2 will be produced by the decomposition of 2, 1, or 5.8 moles of water? Section 9.1 Using Chemical Equations B. Mole-Mole Relationships using Factor Analysis • The mole ratio allows us to convert from moles of one substance in a balanced equation to moles of a second substance in the equation. Use factor analysis and mole ratios to calculate the answers below: CH4(g) + 2O2(g) CO2 + 2H2O(l) 4.2 moles of CH4 reacts with how many moles of O2? 4.2 moles of CH4 will produce how many moles of H2O? (WOC P310 Q8, 10) Section 9.1 Using Chemical Equations Objectives 1. To learn to relate masses of reactants and products in a chemical reaction 2. To perform mass calculations that involve scientific notation Moles Grams Moles Grams Section 9.1 Using Chemical Equations A. Mass Calculations If 9.0 grams of water are decomposed to hydrogen and oxygen, what mass of oxygen gas is formed? (WOC p311 Q15,16) Section 9.1 Using Chemical Equations C. Mass Calculations: Comparing Two Reactions • For equal masses of methane (CH4) and propane (C3H8), which will use more oxygen when fully burned? "Happy Mole Day to You" Chemistry Song Section 9.1 Using Chemical Equations Double Bacon Cheeseburgers and Stoichiometry • 1 Double Bacon Cheeseburger needs 1 bun, 2 patties, 2 slices of cheese, 4 strips of bacon • For 5 Double Bacon Cheeseburgers how many units of each ingredient do I need? • Fill in the final column below with how many complete burgers I can make: Buns Patties Cheese Slices Bacon Strips 2 4 4 8 4 4 8 16 16 30 32 48 How Many Burgers? Section 9.1 Using Chemical Equations Objectives 1. To understand the concept of limiting reactants 2. To learn to recognize the limiting reactant in a reaction 3. To learn to use the limiting reactant to do stoichiometric calculations 4. To learn to calculate percent yield Section 9.1 Using Chemical Equations A. The Concept of Limiting Reactants • Stoichiometric mixture – N2(g) + 3H2(g) 2NH3(g) Section 9.1 Using Chemical Equations A. The Concept of Limiting Reactants • Limiting reactant mixture – N2(g) + 3H2(g) 2NH3(g) Section 9.1 Using Chemical Equations A. The Concept of Limiting Reactants • For a Limiting reactant mixture the number of moles are not balanced to match the reaction equation – N2(g) + 3H2(g) 2NH3(g) – Limiting reactant is the reactant that runs out first – When the limiting reactant is exhausted, then the reaction stops Limiting Reactants Game Section 9.1 Using Chemical Equations B. Calculations Involving a Limiting Reactant Section 9.1 Using Chemical Equations B. Calculations Involving a Limiting Reactant Section 9.1 Using Chemical Equations C. Percent Yield • Theoretical Yield – The maximum amount of a given product that can be formed when the limiting reactant is completely consumed. • The actual yield (amount produced) of a reaction is usually less than the maximum expected (theoretical yield). • Percent Yield – The actual amount of a given product as the percentage of the theoretical yield. Section 9.1 Using Chemical Equations According to her pre-lab theoretical yield calculations a student’s experiment should have produced 1.44g of magnesium oxide. When she weighed her product after reaction, only 1.23g of magnesium oxide was present. What is the student’s percent yield?