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Information Given by the Chemical Equation

Information Given by the Chemical Equation. Balanced equations show the relationship between the relative numbers of reacting molecules and product molecules. 2 CO + O 2  2 CO 2 2 CO molecules react with 1 O 2 molecule to produce 2 CO 2 molecules.

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Information Given by the Chemical Equation

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  1. Information Given by theChemical Equation • Balanced equations show the relationship between the relative numbers of reacting molecules and product molecules. 2 CO + O2 2 CO2 2 CO molecules react with 1 O2 molecule to produce 2 CO2 molecules.

  2. Information Given by theChemical Equation • The information given is relative: 2 CO + O2 2 CO2 Therefore: 200 CO molecules react with 100 O2 molecules to produce 200 CO2 molecules. 2 moles CO molecules react with 1 mole O2 molecules to produce 2 moles CO2 molecules. 12 moles CO molecules react with 6 moles O2 molecules to produce 12 moles CO2 molecules.

  3. Information Given by theChemical Equation • The coefficients in the balanced chemical equation also show the mole ratios of the reactants and products. • Since moles can be converted to masses, we can also determine the mass ratios of the reactants and products.

  4. Information Given by theChemical Equation 2 CO + O2 2 CO2 2 moles CO; 1 mole O2; 2 moles CO2 Since 1 mole of CO = 28.01 g (12.01g + 16g), 1 mole O2 = 32.00 g (2 x 16.00 g), and 1 mole CO2 = 44.01 g (12.01 g + 32 g), then we have the following quantities: 2(28.01) g CO; 1(32.00) g O2; 2(44.01) g CO2

  5. Example #1 Determine the number of moles of carbon monoxide required to react with 3.2 moles oxygen, and determine the moles of carbon dioxide produced.

  6. Example #1 (cont.) • Write the balanced equation: 2 CO + O2 2 CO2 • Use the coefficients to find the mole relationship: 2 moles CO = 1 mol O2 = 2 moles CO2

  7. Example #1 (cont.) • Use dimensional analysis:

  8. Example #2 Consider the following reaction: CH4(g) + 4Cl2(g)  CCl4(l) + 4HCl(g) How many moles of Cl2 would be needed to react with 0.3 mol of CH4?

  9. Example #3 Determine the number of grams of carbon monoxide required to react with 48 g of oxygen, and determine the number of grams of carbon dioxide produced.

  10. Example #3 (cont.) • Write the balanced equation: 2 CO + O2 2 CO2 • Use the coefficients to find the mole relationships: • Determine the molar mass of each:

  11. Example #3 (cont.) • Use the molar mass of the given quantity to convert it to moles. • Use the mole relationships to convert the moles of the given quantity to the moles of the desired quantity:

  12. Example #3 (cont.) • Use the molar mass of the desired quantity to convert the moles to mass:

  13. Example #4 Consider the reaction below: Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g) If you start with 0.056 mol of Zn, how much HCl in grams is needed? How much ZnCl2 in grams can be produced?

  14. Limiting and Excess Reactants • Limiting reactant: a reactant that is completely consumed when a reaction is run to completion • Excess reactant: a reactant that is not completely consumed in a reaction • Theoretical yield: the maximum amount of a product that can be made by the time the limiting reactant is completely consumed

  15. Consider the following reactions: What is the limiting reactant and what is the excess reactant?

  16. Example #5 Determine the number of moles of potassium oxide produced when 8 moles of potassium metal reacts with 3 moles of oxygen.

  17. Example #5 (cont.) • Write the balanced equation: • Use the coefficients to find the mole relationships:

  18. Example #5 (cont.) • Use dimensional analysis to determine the # of moles of reactant A needed to react with reactant B:

  19. Example #5 (cont.) • Compare the calculated # of moles of reactant A to the # of moles of reactant A given. • If the calculated # of moles is greater, then A is the limiting reactant; if the calculated # of moles is less, then A is the excess reactant. • What is the limiting reagent in this reaction?

  20. Example #5 (cont.) • Use the limiting reactant to determine the maximum # of moles of product that can form:

  21. Example #6 Consider the reaction between zinc metal and iodine (I2) to form zinc(II) iodide. If you begin with 50g of Zn and 50 g of I2, what is the limiting reagent? How much zinc(II) iodide in grams can be formed in this reaction?

  22. Example #7 Consider the following unbalanced equation: SF4 + I2O5 IF5 + SO2 Calculate the maximum number of grams of IF5 that can be produced from 10 g of SF4 and 10 g of I2O5.

  23. Actual Yield Theoretical Yield Percent Yield = x 100% Percent Yield • Most reactions do not go to completion. • Theoretical yield: The maximum amount of product that can be made in a reaction • Actual yield: the amount of product actually made in a reaction

  24. Example #8 If in the last problem 9.0 g of IF5 is actually made, what is the percent yield of IF5?

  25. Example #9 Consider the following reaction: N2(g) + 3H2(g)  2NH3(g) When 10g of N2 and 10 g of NH3 react, 8 g of NH3 is formed. What is the percent yield of NH3?

  26. Example #10 Consider the following reaction: CS2(l) + 3O2(g)  CO2(g) + 2SO2(g) When 1 g of CS2 reacts with 1 g of O2, 0.33 g of CO2 results. What is the percent yield of CO2?

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