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Understanding Chemical Quantities and Equations: Moles, Reactants, and Yields

This chapter explores chemical equations and their role in determining the relationships between reactants and products. It discusses how balanced equations convey mole ratios, enabling calculations of moles and masses of involved substances. The concept of limiting and excess reactants is explained, along with methods for calculating theoretical yields and actual yields to determine percent yields. Detailed examples illustrate how to apply dimensional analysis and interpret coefficients in balanced equations for solving chemical quantity problems.

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Understanding Chemical Quantities and Equations: Moles, Reactants, and Yields

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  1. Chapter 9Chemical Quantities

  2. 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 molecules to produce 2 CO2 molecules

  3. Information Given by theChemical Equation • Since the information given is relative: 2 CO + O2 2 CO2 • 200 CO molecules react with 100 O2 molecules to produce 200 CO2 molecules • 2 billion CO molecules react with 1 billion O2 molecules to produce 20 billion 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

  4. Information Given by theChemical Equation • The coefficients in the balanced chemical equation show the molecules and the mole ratio of the reactants and products. • Since moles can be converted to masses, we can determine the mass ratio of the reactants and products as well.

  5. 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, 1 mole O2 = 32.00 g, and 1 mole CO2 = 44.01 g 2(28.01) g CO = 1(32.00) g O2 = 2(44.01) g CO2

  6. 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.

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

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

  9. Steps to Solve Problems • 1) Balance Equation! • 2) Take a ride on the mole road (Dimensional Analysis) • 3) Wanted / Given

  10. Example #2 Determine the number of moles of carbon monoxide required to react with 3.2 moles oxygen, and determine the moles of carbon dioxide produced.

  11. Example #2 (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 • Determine the molar mass of each: 1 mol CO = 28.01 g 1 mol O2 = 32.00 g 1 mol CO2 = 44.01 g

  12. Example #2 (cont.) • Use the molar mass of the given quantity to convert it to moles. • Use the mole relationship to convert the moles of the given quantity to the moles of the desired quantity:

  13. Example #2 (cont.) • Use the molar mass of the desired quantity to convert the moles to mass:

  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. Example #3 Determine the number of moles of carbon dioxide produced when 3.2 moles oxygen reacts with 4.0 moles of carbon monoxide.

  16. Example #3 (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

  17. Example #3 (cont.) • Use dimensional analysis to determine the number of moles of reactant A needed to react with reactant B:

  18. Example #3 (cont.) • Compare the calculated number of moles of reactant A to the number of moles given of reactant A. • If the calculated moles is greater, then A is the limiting reactant; if the calculated moles is less, then A is the excess reactant. • The calculated moles of CO (6.4 moles) is greater than the given 4.0 moles; therefore CO is the limiting reactant.

  19. Example #3 (cont.) • Use the limiting reactant to determine the moles of product:

  20. Example #4 Determine the mass of carbon dioxide produced when 48.0 g of oxygen reacts with 56.0 g of carbon monoxide.

  21. Example #4 (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 • Determine the molar mass of each: 1 mol CO = 28.01 g 1 mol O2 = 32.00 g 1 mol CO2 = 44.01 g

  22. Example #4 (cont.) • Determine the moles of each reactant:

  23. Example #4 (cont.) • Determine the number of moles of reactant A needed to react with reactant B:

  24. Example #4 (cont.) • Compare the calculated number of moles of reactant A to the number of moles given of reactant A. • If the calculated moles is greater, then A is the limiting reactant; if the calculated moles is less, then A is the excess reactant. • The calculated moles of O2 (1.00 moles) is less than the given 1.50 moles; therefore O2 is the excess reactant.

  25. Example #4 (cont.) • Use the limiting reactant to determine the moles of product, then the mass of product:

  26. Actual Yield Theoretical Yield Percent Yield = x 100% Percent Yield • Most reactions do not go to completion. • Actual yield: the amount of product made in an experiment • Percent yield: the percentage of the theoretical yield that is actually made

  27. Example #4a If in the last problem 72.0 g of carbon dioxide is actually made, what is the percentage yield?

  28. Example #4a (cont.) • Divide the actual yield by the theoretical yield, then multiply by 100% • The actual yield of CO2 is 72.0 g • The theoretical yield of CO2 is 88.0g

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