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Chapter 9 Chemical Quantities. 9 | 2 Information Given by the Chemical Equation Balanced equations show the relationship between the relative numbers.

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Presentation on theme: "Chapter 9 Chemical Quantities. 9 | 2 Information Given by the Chemical Equation Balanced equations show the relationship between the relative numbers."— Presentation transcript:

1 Chapter 9 Chemical Quantities

2 9 | 2 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 molecules to produce 2 CO 2 molecules

3 9 | 3 Since the information given is relative: 2 CO + O 2  2 CO CO molecules react with 100 O 2 molecules to produce 200 CO 2 molecules 2.2 billion CO molecules react with 1 billion O 2 molecules to produce 20 billion CO 2 molecules 3.2 moles CO molecules react with 1 mole O 2 molecules to produce 2 moles CO 2 molecules 4.12 moles CO molecules react with 6 moles O 2 molecules to produce 12 moles CO 2 molecules Information Given by the Chemical Equation

4 9 | 4 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. Information Given by the Chemical Equation

5 9 | 5 2 CO + O 2  2 CO 2 2 moles CO = 1 mole O 2 = 2 moles CO 2 Since 1 mole of CO = g, 1 mole O 2 = g, and 1 mole CO 2 = g 2(28.01) g CO = 1(32.00) g O 2 = 2(44.01) g CO 2 Information Given by the Chemical Equation

6 9 | 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 9 | 7 Example #1 (cont.) Write the balanced equation: 2 CO + O 2  2 CO 2 Use the coefficients to find the mole relationship: (Wanted/Given) 2 moles CO = 1 mol O 2 = 2 moles CO 2

8 9 | 8 Use dimensional analysis: Example #1 (cont.)

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

10 9 | 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 9 | 11 Example #2 (cont.) Write the balanced equation: 2 CO + O 2  2 CO 2 Use the coefficients to find the mole relationship: 2 moles CO = 1 mol O 2 = 2 moles CO 2 Determine the molar mass of each: 1 mol CO = g 1 mol O 2 = g 1 mol CO 2 = g

12 9 | 12 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: Example #2 (cont.)

13 9 | 13 Use the molar mass of the desired quantity to convert the moles to mass: Example #2 (cont.)

14 9 | 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 9 | 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 9 | 16 Example #3 (cont.) Write the balanced equation: 2 CO + O 2  2 CO 2 Use the coefficients to find the mole relationship: 2 moles CO = 1 mol O 2 = 2 moles CO 2

17 9 | 17 Use dimensional analysis to determine the number of moles of reactant A needed to react with reactant B: Example #3 (cont.)

18 9 | 18 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. Example #3 (cont.)

19 9 | 19 Use the limiting reactant to determine the moles of product: Example #3 (cont.)

20 9 | 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 9 | 21 Example #4 (cont.) Write the balanced equation: 2 CO + O 2  2 CO 2 Use the coefficients to find the mole relationship: 2 moles CO = 1 mol O 2 = 2 moles CO 2 Determine the molar mass of each: 1 mol CO = g 1 mol O 2 = g 1 mol CO 2 = g

22 9 | 22 Determine the moles of each reactant: Example #4 (cont.)

23 9 | 23 Determine the number of moles of reactant A needed to react with reactant B: Example #4 (cont.)

24 9 | 24 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 O 2 (1.00 moles) is less than the given 1.50 moles; therefore O 2 is the excess reactant. Example #4 (cont.)

25 9 | 25 Use the limiting reactant to determine the moles of product, then the mass of product: Example #4 (cont.)

26 9 | 26 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 Percent Yield = Actual Yield Theoretical Yield x 100%

27 9 | 27 Example #4a If in the last problem 72.0 g of carbon dioxide is actually made, what is the percentage yield ?

28 9 | 28 Example #4a (cont.) Divide the actual yield by the theoretical yield, then multiply by 100% –The actual yield of CO 2 is 72.0 g –The theoretical yield of CO 2 is 88.0g


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