Chapter 3: Chemical Reactions and the Earth

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Chapter 3 Chemical Reactions and the Earths
Composition

• Problems 3.1-3.3, 3.5, 3.11-3.86, 3.95-3.115,
  3.119-3.120, 3.122, 3.125-3.128, 3.132, 3.134,
  3.136-3.138-3.141

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3.2 The Mole

• Stoichiometry (STOY-key-OM-e-tree) quantitative
  study of reactants and products in a chemical
  reaction
• Interpreting a Chemical Equation
• H2 (g) Cl2 (g) ? 2 HCl (g)
• 1 molecule 1 molecule 2 molecules
• It follows that any multiples of these
  coefficients will be in same ratio!
• 2 H2 (g) O2 (g) ?
  2 H2O(g)
• ?1000 _____ molecule(s) _____
  molecule(s) _____ molecule(s)
• ?N _____ molecule(s) _____
  molecule(s) _____ molecule(s)
• Since N Avogadros 6.022?1023 molecules
  1 mole
• 2 H2 (g) O2 (g) ? 2
  H2O(g)
• _____ mole(s) _____ mole(s)
  _____ mole(s)
• Thus, the coefficients in a chemical equation
  give the mole ratios of reactants and products.

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Example Problem

• Consider the following
• 2 C2H6 (g) 7 O2 (g) ? 4 CO2 (g) 6
  H2O (g)
•  
• 1. How many moles of O2 will react with 2.50
  moles of C2H6?
• 2. How many moles of CO2 form when 3.50 moles of
  O2 completely react?
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Stoichiometric Calculations and the Carbon cycle

• Mass-Mass Stoichiometry Problems
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• Example 1 Photosynthesis is the process of
  energy from sunlight being used to convert carbon
  dioxide into organic compounds, especially sugars
  like glucose, C6H12O6
• 6 CO2(g) 6 H2O(g) ? C6H12O6(aq)
  6 O2(g)
• What mass (in g) of glucose is produced via
  photosynthesis when 25.0 kg of carbon dioxide
  react with excess steam?
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6 CO2(g) 6 H2O(g) ? C6H12O6(aq)
6 O2(g)

• What mass (in g) of glucose is produced via
  photosynthesis when 25.0 kg of carbon dioxide
  react with excess steam?

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Stoichiometric Calculations and the Carbon cycle

• Example 2 In biological systems, the reverse
  reaction,
•  
• C6H12O6(aq) 6 O2(g) ? 6 CO2(g)
  6 H2O(g)
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• is called respiration and is the major source of
  energy for all livings things.
• a. Calculate the mass (in g) of carbon dioxide
  produced when 5.00 lb. of glucose reacts
  completely. (1 lb. 453.6 g)
•  
• b. How many pounds of carbon dioxide are produced
  in problem a. above?

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C6H12O6(aq) 6 O2(g) ? 6 CO2(g)
6 H2O(g)

• Calculate the mass (in g) of carbon dioxide
  produced when 5.00 lb. of glucose reacts
  completely.
•  
• b. How many pounds of carbon dioxide are produced
  in problem a. above?

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3.4 Combustion Reactions

• CxHy O2(g) ? CO2(g) H2O(g)
• CxHyOz O2(g) ? CO2(g) H2O(g)
•  
• Hydrocarbons (compounds with only C and H) and
  hydrocarbon derivatives (compounds with only C, H
  and O) burn in O2 to produce CO2 gas and steam,
  H2O(g).
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Combustion Reactions

• Example 1 Many home barbecues are fueled with
  propane gas (C3H8). Write the balanced equation
  for the combustion of propane, then calculate the
  mass (in kg) of carbon dioxide produced upon
  complete combustion of liquid propane from a 5.0
  gal tank.
• (The density of liquid propane at 60F is about
  4.2 lbs. per gallon, and 1 lb. 453.6 g)

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Combustion Reactions

• Example 2 Everclear is a brand of grain alcohol
  that can be as high as 190 proof (or 95 ethanol,
  C2H5OH, by volume). Calculate the mass of carbon
  dioxide produced upon complete combustion of the
  ethanol in a 750 mL bottle of Everclear. Write
  the balanced chemical equation for the combustion
  of ethanol. (The density of this Everclear is
  0.80 g/mL.)

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3.9 Limiting Reactants and Percent Yield

• In practice, reactants will not always be present
  in the exact amounts necessary to be converted
  completely into products.
• Some reactants (usually the more expensive) are
  only present in a limited supply, so these are
  almost always completely used up
• limiting reactant (or limiting reagent) since
  it limits the amount of product made
• Some reactants (usually the less expensive) are
  present in larger amounts and are never
  completely used up ? reactant(s) in excess

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Guidelines for solving Limiting Reactant Problems

• Calculate the mass or the of moles of the 2nd
  reactant needed to completely react with the 1st
  reactant.
• If the moles needed is greater than the number of
  moles present for the 2nd reactant
• That 2nd reactant will run out before the 1st
  reactant.
• The 2nd reactant the limiting reactant, and the
  1st reactant is in excess.
• If the moles needed is less than the number of
  moles present for the reactant,
• The 1st reactant the limiting reactant, and the
  2nd reactant is in excess.
• Use the amount of the limiting reactant present
  to solve for the mass or of moles of product
  that can be made.

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Limiting Reactant Problems

• Consider the reaction to produce ammonia
• N2(g) 3 H2(g) ? 2 NH3(g)
• Example 1 a) If 50.0 g of N2 react with 10.0 g
  of H2, what mass of ammonia is produced?
• b) The limiting reactant is _______ and the
  excess reactant is _________.
• c) What mass of the reactant in excess remains
  after the reaction?

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Calculating Percent Yield

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Calculating Percent Yield

• Example 1 N2(g) 3 H2(g) ? 2 NH3(g)
• a. For the reaction of 50.0 g of N2 with 10.0 g
  of H2, the theoretical yield of ammonia was
  determined to be what?
• theoretical yield
• If 49.6 g of ammonia were actually produced,
  calculate the percent yield for the reaction.
• percent yield

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Calculating Percent Yield

• Example 2
• Consider the following reaction
• 2 KClO3(s) ? 2 KCl(s) 3 O2(g)
•  
• What is the percent yield if 50.0 g of KClO3
  decomposes to produce 16.4 g of oxygen gas?

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Calculating Percent Yield

• Example 3
• Consider the following reaction
•  
• 3 Na2CrO4 (aq) 2 AlCl3 (aq) ?
  Al2(CrO4)3 (s) 6 NaCl (aq)
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• a. What mass of precipitate is produced when 50.0
  g of sodium chromate react with 50.0 g of
  aluminum chloride? Which is the limiting reactant
  and which is the reactant in excess?

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Calculating Percent Yield

• Example 3
• Consider the following reaction
•  
• 3 Na2CrO4(aq) 2 AlCl3(aq) ?
  Al2(CrO4)3(s) 6 NaCl(aq).
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• b. What mass of the reactant in excess remains
  after the reaction?
• c. What is the percent yield if 4.32 g of
  precipitate is actually produced?

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Calculating Percent Yield

• Example 4
• Calculate the mass of methane (CH4) that must
  react to produce 10.0 kg of carbon dioxide if the
  percent yield for the reaction is 88.8.

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Calculating Percent Yield

• Example 5
• Consider the thermal decomposition of N2O5
• 2 N2O5(g) 4 NO2(g) O2(g)
• If the percent yield for the reaction is 96.8,
  and the density of oxygen gas is 1.31 g/L,
  calculate the mass of N2O5 required to produce
  50.0 L of oxygen gas.