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Chapter 4:

Chapter 4:. Balancing Equations and Stoichiometry. Key Terms & Concepts. Stoichiometry Chemical Equations reactants and products balancing chemical equations Chemical Calculations Limiting Reactant Theoretical and Percent Yield. Stoichiometry.

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Chapter 4:

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  1. Chapter 4: Balancing Equations and Stoichiometry

  2. Key Terms & Concepts • Stoichiometry • Chemical Equations • reactants and products • balancing chemical equations • Chemical Calculations • Limiting Reactant • Theoretical and Percent Yield

  3. Stoichiometry • Stoichiometry is the study of the quantitative nature of chemical formulas and chemical reactions. • Stoichiometry is one the the most essential tools in chemistry • It allows to quantify everything from global warming to drug manufacturing

  4. Chemical Equations • Chemical reactions are represented in a concise manner by chemical equation • For example, when H2 burns in O2, H2O is formed. • The chemical equation for this reaction is: 2 H2 + O2 2H2O

  5. Chemical Equations 2 H2 + O2 2 H2O • The compounds on the left of the arrow are called “reactants” • The compounds on the right of the arrow are called “products” • H2 and O2 are reactants, H2O is the product

  6. Chemical Equations 2 H2 + O2 2 H2O • Notice that the number of atoms or each element is equal on both sides of the equation • 4 H, 2 0 • All chemical equations must meet this requirement • Chemical equations must be balanced!! • We balance equations by changing coefficients, not chemical formulas

  7. Chemical Equations

  8. Chemical Equations • Consider the following chemical equation CH4 + O2 CO2 + H2O unbalanced • Start with elements that only appear in one compound on either side of the equation • C and H are only in one compound on either side • C is balanced

  9. Chemical Equations CH4 + O2 CO2 + H2O unbalanced • 4 H’s in reactants, 2 H’s in products • Put coefficient of 2 in front of H2O • 4 H’s in reactants, 4 H’s in products CH4 + O2 CO2 + 2 H2O unbalanced

  10. Chemical Equations CH4 + O2 CO2 + 2 H2O unbalanced • 2 O’s in reactants, 4 O’s in products • Put coefficient of 2 in front of O2 • 4 O’s in reactants, 4 O’s in products CH4 + 2 O2 CO2 + 2 H2O balanced

  11. Chemical Equations

  12. Chemical Equations • Consider this equation: C3H8 + O2  CO2 + H2O 3C + 8H + 2O  1C + 2H + 3O

  13. Chemical Equations • Balance C and H C3H8 + O2  3CO2 + 4H2O 3C + 8H + 2O  3C + 8H + 10O

  14. Chemical Equations • Balance O C3H8 + 5O2  3CO2 + 4H2O 3C + 8H + 10O  3C + 8H + 10O

  15. Chemical Equations • Example 4.1 Balance the following chemical equations (1) Mg + HCl  MgCl2 + H2 (2) K + H2O  KOH + H2 (3) CaCl2 + Na3PO4  Ca3(PO4)2 + NaCl (4) NaN3  Na + N2 (5) C8H18 + O2  CO2 + H2O

  16. Chemical Equations • Example 4.1 Balance the following chemical equations (1) Mg + 2 HCl  MgCl2 + H2 (2) 2 K + 2 H2O  2 KOH + H2 (3) 3 CaCl2 + 2 Na3PO4  Ca3(PO4)2 + 6 NaCl (4) 2 NaN3  2 Na + 3 N2 (5) 2 C8H18 + 25 O2  16 CO2 + 18 H2O

  17. Chemical Equations • Example 4.2 Write a balanced chemical equation for the following reactions • ammonium nitrate decomposes to nitrogen gas, oxygen gas, and water • iron reacts with oxygen gas and water to form iron(II) hydroxide • ammonia reacts with oxygen gas to produce nitrogen monoxide and water

  18. Chemical Equations • Example 4.2 Write a balanced chemical equation for the following reactions • 2 NH4NO3  2 N2 + O2 + 4 H2O • 2 Fe + O2 + 2 H2O  2 Fe(OH)2 • 4NH3 + 5 O2  4 NO + 6 H2O

  19. Chemical Calculations 2 H2 + O2 2 H2O 2 molecules 1 molecule 2 molecules 2(6.022x1023) molecules 6.022x1023 molecules 2(6.022x1023) molecules 2 mol 1 mol 2 mol • Stoichiometric coefficients can be interpreted as either number of molecules or number of moles.

  20. Chemical Calculations • Example 4.3 How many moles of water can be produced from 5.25 mol O2?

  21. Chemical Calculations • Example 4.4 How many moles of oxygen are required to completely react with 8.50 moles of butane, C4H10?

  22. Chemical Calculations • We can’t directly measure moles. We can measure mass. • We can use the stoichiometric coefficients of a reaction to determine the mass relationships. • However, we must always convert mass to moles. • We cannot directly compare the masses of reactants and products. • We can only compare the moles of reactants and products.

  23. Chemical Calculations • The general scheme is:

  24. Chemical Calculations • Example 4.5 Geranyl formate is used as a synthetic rose essence in cosmetics. The compound is prepared from formic acid and geraniol: HCOOH + C10H18O  C11H18O2 + H2O A chemist needs to make some geranyl formate for a batch of perfume. How many grams of geranyl formate can a chemist make from 375g of geraniol?

  25. Chemical Calculations

  26. Chemical Calculations • Example 4.6 Solid lithium hydroxide is used in space vehicles to remove exhaled carbon dioxide. The lithium hydroxide reacts with gaseous carbon dioxide to form solid lithium carbonate and liquid water. How many grams of carbon dioxide can be absorbed by each 1.00 g of lithium hydroxide? 2 LiOH (s) + CO2 (g)  Li2CO3 (g) + H2O (l)

  27. Chemical Calculations • Example 4.6 2 LiOH (s) + CO2 (g)  Li2CO3 (g) + H2O (l)

  28. Theoretical and Percent Yield • The amount of product that can be produced from a given amount of reactants is the theoretical yield. • However, no reaction goes to actual completion. The amount of products that is actually produced from a given amount of reactants is the actual yield. • Some reactants may not react • Reactants may react in an undesired way (side reactions) • May be difficult to remove products from pot

  29. Theoretical and Percent Yield • The extent of the desired reaction is typically reported as the percent yield.

  30. Theoretical and Percent Yield • Example 4.7 Look back at Example 4.5. If the chemist starts with 375g of geraniol and collects 417g of purified product, what is the percent yield of the synthesis?

  31. Theoretical and Percent Yield • Example 4.8 25.0 g of sodium metal is burned in an excess of chlorine gas. What is the theoretical yield of sodium chloride? If 54.8 g of sodium chloride is actually produced, what is the percent yield of the reaction? 2 Na + Cl2 2 NaCl

  32. Theoretical and Percent Yield

  33. Theoretical and Percent Yield

  34. Theoretical and Percent Yield • Example 4.9 Titanium is a strong, lightweight, corrosion-resistant metal that is used in aeronautics and bicycle frames. It is prepared by the reaction of titanium (IV) chloride with molten magnesium between 950C and 1150C. TiCl4(g) + 2 Mg(l)  Ti(s) + 2 MgCl2(l) In a certain industrial process 3.54x107 g of TiCl4 are reacted with 1.13x107 g of Mg. (a)Calculate the theoretical yield of Ti in grams. (b)Calculate the percent yield if 7.91x106 g of Ti are actually produced.

  35. Theoretical and Percent Yield

  36. Theoretical and Percent Yield

  37. Limiting Reactants • Most reactions do not occur with stoichiometric equivalent amounts of each reactant. • One reactant is used up first • This reactant is the limiting reactant because it limits the amount of products that can be formed

  38. Limiting Reactants • Consider the “ham sandwich” example • one sandwich is made from one slice of ham, one slice of cheese and two slices of bread • How many ham sandwiches can be made from six slices of ham, seven slices of cheese and 14 slices of bread? • What is the limiting reactant?

  39. Limiting Reactants

  40. Limiting Reactants • If a problem gives specific amounts of two or more reactants it is a limiting reactant problem. • Determine the amount of product that can be formed from each reactant • The reactant which produces the smallest amount of product is the limiting reactant • The remaining reactants are said to be in excess

  41. Limiting Reactants • Example 4.10 How many moles of water can be formed when 10.0 moles of H2 reacts with 4.50 moles of O2? What is the limiting reactant?

  42. Limiting Reactants • Example 4.10 2 H2 + O2 2 H2O

  43. Limiting Reactants • Example 4.11 Solutions of sulfuric acid and lead (II) acetate react to form solid lead (II) sulfate and aqueous acetic acid. If 15.0 g of sulfuric acid and 15.0 g of lead (II) acetate are mixed, calculate the number of grams of lead (II) sulfate that can be produced. Also calculate the number of grams of the excess reagent remaining after the reaction is completed.

  44. Limiting Reactants H2SO4 + Pb(CH3COO)2 PbSO4 + 2 CH3COOH

  45. Limiting Reactants

  46. End-of-Chapter Exercises • Suggested End-of-Chapter Exercises 5, 9, 10, 14, 17, 20, 22, 28

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