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Stoichiometry

Stoichiometry. Chapter 12. 12. 1 What is stoichiometry?. Stoichiometry is the study of quantitative relationships between the amounts of reactants used and products formed by a chemical reaction.

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Stoichiometry

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  1. Stoichiometry Chapter 12

  2. 12. 1 What is stoichiometry? • Stoichiometry is the study of quantitative relationships between the amounts of reactants used and products formed by a chemical reaction. • It is based on the law of conservation of mass, which says that matter is not created or destroyed.

  3. Balanced Chemical Equations • The coefficients in a balanced equation can tell you two things: • How many individual particles are interacting in the chemical reaction. • How many moles of particles interact to form the reaction. • The coefficients do not tell you directly the mass of reactants and products in the reaction.

  4. Problems • Interpret each balanced equation in terms of particles, moles, and mass. Show that the law of conservation of mass is observed. • 2H2O2 (l) O2 (g) + 2H2O (l) • H2CO3 (aq)  H2O (l) + CO2 (g) • 4HCl (aq) + O2(g)  2H2O (l) + 2Cl2 (g)

  5. Mole Ratio • Since coefficients show relationships in the moles of reactants and products, they can be used to write mole ratios. • A mole ratio is a ratio between the number of moles of any two substances in a balanced chemical equation.

  6. Problems • Determine all the mole ratios for the following balanced chemical equations. • N2 (g) + O2 (g)  2NO (g) • 4NH3 (aq) + 5O2 (g)  4NO (g) +6H2O (l) • 4HCl (aq) + O2 (g)  2H2O (l) + 2Cl2 (g)

  7. Stiochiometric Calculations Section 12.2

  8. Stoichiometric Calculations • Three basic stoichiometric calculations: • Mole-to-mole conversions • Mole-to-mass conversions • Mass-to-mass conversions

  9. Steps for Calculations • The first thing you need, before you can do any stoichiometric calculations, is to have a balanced chemical equation. • Remember, when you are converting from one substance to another you must compare the quantities using the units moles. • Moles are the only basis of comparison between two different substances.

  10. Mole-to-Mole Conversions • How can you determine the number of moles of table salt (NaCl) produced from 0.02 moles of chlorine (Cl2)? • First, begin with the balanced chemical equation. • 2Na (s) + Cl2 (g)  2NaCl (s) • Then use the mole ratio to convert the known moles of chlorine to the number of moles of table salt.

  11. Example • A piece of magnesium burns in the presence of oxygen, forming magnesium oxide (MgO). How many moles of oxygen are needed to produce 12 moles of magnesium oxide? • First, write the equation. • 2Mg (s) + O2 (g) 2MgO (s) • 1 mole oxygen = 2 molmagnesiumoxide • Set up a dimensional analysis problem to solve.

  12. Problem • The carbon dioxide exhaled by astronauts can be removed from a spacecraft by reacting it with lithium hydroxide (LiOH). The reaction is as follows: CO2 (g) + 2LiOH (s) Li2CO3 (s) + H2O (l). An average person exhales about 20 moles of CO2 per day. How many moles of LiOH would be required to maintain two astronauts in a spacecraft for three days?

  13. Problem • Balance the following equation and answer the question below. KClO3 (s)  KCl (s) + O2 (g) • How many moles of O2 are produced from 10 moles of KClO3? • How many moles of KCl are produced using 3 moles of KClO3? • How many moles of KClO3 are needed to produce 50 moles of O2?

  14. Mole-to-Mass conversion • Remember, all mass to mole conversions are completed by using the molar mass (the mass of 1 mole of the substance). • Set up a dimensional analysis problem to solve.

  15. Problem • Calculate the mass of sodium chloride (NaCl) produced when 5.50 moles of sodium react in excess chlorine gas. • How many grams of chlorine gas must be reacted with excess sodium iodide (NaI) to produce 6.00 moles of sodium chloride? • Balance the equation: NaI (aq) + Cl2 (g)  NaCl (aq) + I2 (s) • Perform the calculation

  16. Mass-to-Mass Conversions • In order to compare two different quantities remember that you must use the mole ratio. • So, first convert the mass of the first substance to moles. • Then convert the moles of the first substance to moles of the second substance. • Then convert the moles of the second to mass of the second substance.

  17. Example • How many grams of sodium hydroxide (NaOH) are needed to completely react with 50.0 grams of sulfuric acid (H2SO4) to form sodium sulfate (Na2SO4) and water?

  18. Problem • Balance each equation and solve the problem. • If 40.0 g of magnesium reacts with excess hydrochloric acid (HCl), how many grams of magnesium chloride (MgCl2) are produced? Mg (s) + HCl (aq)  MgCl2 (aq) + H2 (g) • Determine the mass of copper needed to react completely with a solution containing 12.0 g of silver nitrate (AgNO3)? Cu (s) + AgNO3 (aq)  Cu(NO3)2 (aq) + Ag (s) • How many grams of hydrogen chloride (HCl) are produced when 15.0 g of sodium chloride (NaCl) reacts with excess sulfuric acid (H2SO4)? NaCl (aq) + H2SO4 (aq) Na2SO4 (s) + HCl (g)

  19. Section 12.3 Limiting reactants

  20. Limiting Reactants • Rarely are the reactants in a chemical reaction present in the exact mole ratio shown in the balanced equation. • When one reactant is used completely up first, it is called the limiting reactant. • The limiting reactant limits the reaction and thus determines how much product forms. • The left-over reactants are called the excess reactants.

  21. Limiting Reactant • To find out which reactant is the limiting reactant, convert each of the reactants to moles of product. • The reactant that produces the smallest amount of product is the limiting reactant. • The other reactants are excess reactants.

  22. Excess Reactant • Once the limiting reactant has been found, you can find the amount of excess that is left-over after the reaction is finished. • To do this, convert the limiting reactant to mass of the excess reactant. • Then subtract the mass of the excess from the initial amount of the excess reactant.

  23. Example • In the reaction below, 40.0 g of sodium hydroxide (NaOH) reacts with 60.0 g of sulfuric acid (H2SO4). 2NaOH (aq) + H2SO4 (aq)  Na2SO4 (aq) + 2H2O (g) • Which reactant is the limiting reactant? • What mass of Na2SO4 can be produced using the given quantities of the reactants?

  24. Movie • Limiting Reactants

  25. Problems • Ammonia (NH3) is one of the most common chemicals produced in the United States. It is used to make fertilizer and other products. Ammonia is produced by the following chemical reaction. N2 (g) + 3H2 (g)  2NH3 (g) • If you have 1.00 X 103 g of N2 and 2.50 X 103 g of H2, which is the limiting reactant in the reaction? • How many grams of ammonia can be produced from the amount of limiting reactant available? • Calculate the mass of excess reactant that remains after the reaction is complete.

  26. Problem • Aluminum reacts with chlorine to produce aluminum chloride. • Balance the equation: Al (s) + Cl2 (g)  AlCl3 (s) • If you begin with 3.2 g of aluminum and 5,4 g of chlorine, which is the limiting reactant? • How many grams of aluminum chloride can be produced from the amount of limiting reactant available? • Calculate the mass of excess reactant the remains after the reaction is complete.

  27. Section 12.4 Percent yield

  28. Yield • Most reactions do not produce the predicted amount of product. • Many reactions stop before all the reactants are used up, so less product is formed than expected. • Also, sometimes products other than those expected are formed, and so they reduce the amount of the desired product.

  29. Theoretical Yield • The theoretical yield is the maximum amount of product that can be produced from a given amount of reactant under ideal circumstances. • It is the amount that you calculate using stoichiometry.

  30. Actual Yield • The actual yield is the amount of product that is actually produced in an experiment. • It is what you get when you do the lab.

  31. Percent Yield • The percent yield is the ratio of the actual yield to the theoretical yield expressed as a percent. • Percent yield = Actual yield * 100% Theoretical yield • The percent yield tells you how efficient a chemical reaction is in producing the desired product.

  32. Example • Aspirin (C9H8O4) can be made from salicylic acid (C4H6O3). Suppose you mix 13.2 g of salicylic acid with an excess of acetic anhydride and obtain 5.9 g of aspirin and some water. Calculate the % yield of aspirin. 2C7H6O3 (s) + C4H6O3 (l) 2C9H8O4 (s) + H2O (l)

  33. Problem 11. Calculate the percent yield for each chemical reaction based on the data provided. • Theo. Yield: 25 g; act. Yield: 20 g • Theo. Yield: 55g; act. Yield: 42g • Theo. Yield: 5.2 g; act. Yield: 4.9 g

  34. Problem • In an experiment, 10.0 g of magnesium reacted with excess hydrochloric acid forming magnesium chloride. Mg (s) + 2HCl (aq)  MgCl2 (aq) + H2 (g) At the completion of the reaction, 29.5 g of magnesium chloride was produced. Calculate the theoretical yield and the percent yield.

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