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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. - PowerPoint PPT Presentation
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Copyright © Houghton Mifflin Company. All rights reserved. 9 | 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.
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• 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.
Information Given by theChemical Equation
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• 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.
Information Given by theChemical Equation
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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
Information Given by theChemical Equation
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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.
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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
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• Use dimensional analysis:
Example #1 (cont.)
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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?
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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.
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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:
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• 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:
Example #3 (cont.)
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• Use the molar mass of the desired quantity to convert the moles to mass:
Example #3 (cont.)
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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?
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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
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+
+
+
Consider the following reactions:
What is the limiting reactant and what is the excess reactant?
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Example #5
Determine the number of moles of potassium oxide produced when 8 moles of potassium metal reacts with 3 moles of oxygen.
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Example #5 (cont.)
• Write the balanced equation:
• Use the coefficients to find the mole relationships:
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• Use dimensional analysis to determine the # of moles of reactant A needed to react with reactant B:
Example #5 (cont.)
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• 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?
Example #5 (cont.)
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• Use the limiting reactant to determine the maximum # of moles of product that can form:
Example #5 (cont.)
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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?
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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.
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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
Percent Yield = Actual Yield
Theoretical Yieldx 100%
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Example #8
If in the last problem 9.0 g of IF5 is actually made, what is the percent yield of IF5?
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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?
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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?