Stoichiometry Chapter 10

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Stoichiometry Chapter 10

• Mole Relationships
• Types of Stoichiometry Problems
• Mass Relationships in Chemical Reactions
• Limiting Reactants
• Theoretical, Actual and Percent Yields

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Stoichiometry

• Stoichiometry
• The study of quantitative relationships between
  substances undergoing chemical changes.
• Law of Conservation of Matter
• In chemical reactions, the quantity of matter
  does not change.
• The total mass of the products must equal that of
  the reactants.

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Mass Relationshipsin Chemical Reactions

• Stoichiometry - The calculation of quantities of
  reactants and products in a chemical reaction.

H2 O2 ? H2O
2 H2 O2 ? 2 H2O
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Mole calculations

• The balanced equation shows the mole ratio
  between reactants and products.
• 2C2H6 7O2 4CO2 6H2O
• For each chemical, you can determine the
• moles of each reactant consumed
• moles of each product made
• There are many types of stoichiometry problems.
• All involve mole-mole relationships.

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Mole-Mole Relationships

• From Example Exercise 10.3
• Carbon monoxide is produced in a blast furnace by
  passing oxygen gas over hot coal (elemental
  carbon).
• How many moles of oxygen react with 2.50 moles of
  carbon?

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Mole-Mole Relationships

• Self-Test Exercise
• Iron is produced from iron ore by passing carbon
  monoxide gas through molten iron(III) oxide.
• How many moles of carbon monoxide react with 2.50
  moles of iron(III) oxide?
• How many moles of iron would be produced?

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Mass Calculations

• We dont directly weigh out molar quantities.
• Rememberwe dont have a mole-o-meter in the lab
• We can use measured masses like kilograms, grams
  or milligrams.
• The molar masses and the chemical equations allow
  us to use either mass or molar quantities.

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Mass-Mass

• Suppose we have a problem like Example Exercise
  10.3 or even the Self-Test Exercise
• If we were told the mass of one of the
  substances, would we be able to determine the
  number of moles?
• Sure!
• We would then proceed as we did in the previous
  examples.
• Once we obtained the moles of the substance asked
  about in the problem, could we express that
  quantity in grams?
• Sure!
• We have now talked through a mass-mass
  stoichiometry problem
• Lets try one

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Mass Calculations

• How many grams of hydrogen will be produced if
  10.0 grams of calcium is added to an excess of
  hydrochloric acid?
• 2HCl Ca ______gt CaCl2 H2
• Note
• We produce one H2 for each calcium.
• There is an excess of HCl so we have all we need.

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Mass Calculations

• 2HCl Ca ____gt CaCl2 H2
• First - Determine the number of moles of
  calcium available for the reaction.

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Mass Calculations

• 2HCl Ca _____gt CaCl2 H2
• According to the chemical equation, we get one
  mole of H2 for each mole of Ca.
• So we will make 0.25 moles of H2.
• grams H2 produced moles x molar mass H2
• 0.25 mol x 2.016 g/mol
• 0.504 grams

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Mass Calculations

• OK, so how many grams of CaCl2 were made?
• 2HCl Ca _____gt CaCl2 H2
• We would also make 0.25 moles of CaCl2.
• g CaCl2 0.25 mol x FM CaCl2
• 0.25 mol x 110.98 g / mol CaCl2
• 27.75 g CaCl2

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Putting It All Together

• 2 HCl Ca ? CaCl2 H2
• How many grams of H2 can be produced from 10.0
  grams of calcium and excess hydrochloric acid?
• How many grams of CaCl2 can be produced from 10.0
  grams of calcium and excess hydrochloric acid?

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Another Mass-Mass Example

• Solid iron(III) oxide reacts with aluminum metal.
• Suppose we have 15.0 grams of iron(III) oxide.
  How much aluminum, in grams, do we need to react
  with it?

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Mass-Mass Example (cont)

• Using the quantities from the previous
  examplehow many grams of each product could we
  obtain?

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Limiting reactant

• Limiting reactant - the material that is in the
  shortest supply based on a balanced chemical
  equation.

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Limiting reactant animation
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Limiting Reactants

• We deal with limiting reactants every day.
• Have you ever been cooking and run out of some
  ingredient you need?
• You have enough of all the ingredients except
  one.
• What happens to your cooking??
• It has to stop!

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Limiting Reactant Example

• Youre making PBJ sandwiches.
• For each sandwich you need 2 tbsp. peanut butter,
  1 tbsp. jelly, and two slices of bread.
• You have 22 slices of bread, 15 tbsp. jelly, and
  18 tbsp. peanut butter
• How many sandwiches can you make?

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Example

• For the following reaction, which is limiting if
  you have 5.0 g of hydrogen and 10. g oxygen?
• Balanced Chemical Reaction
• 2H2 O2 ________gt 2H2O
• You need 2 moles of H2 for each mole of O2.
• Moles of H2 5.0 g 2.5 mol
• Moles of O2 10.g 0.31 mol

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Example (cont)

• 2 H2 (g) O2 (g) ? 2 H2O (g)
• You have 0.25 mol H2 and 0.31 mol O2
• How many moles of water can you make from each?
• O2 is the limiting reactant, because you can only
  make 0.62 mol H2O.
• It doesnt matter that you have enough H2 to make
  2.5 mol H2O, you will run out of O2 once youve
  made 0.62 mol H2O.
• Same concept as in our practical exampleIt
  doesnt matter how much bread and jelly we have,
  there is only enough peanut butter to make nine
  sandwiches.

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Example (cont)

• Lets add another wrinkle
• Suppose the question asks how many GRAMS of water
  we can make.
• We just take the moles we already calculated and
  convert to grams.

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Putting it all together
Graphically 2 H2 (g) O2 (g) ? 2
H2O (g)
Grams O2
Grams H2O Moles O2
Moles H2O
Molar mass of O2
Molar mass of H2O
Mole ratio from balanced equation
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The Reactant in Excess

• We produced 11.3 grams of water.
• We consumed all 10.0 grams of oxygen in the
  process.
• We did not use all 5.00 grams of hydrogen.
• How much of the hydrogen did we use?
• How much hydrogen did we have left over?

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The Reactant in Excess

• How much hydrogen does it take to produce 11.3
  grams of water, based upon the balanced equation
  below?
• 2 H2 O2 ? 2 H2O
• We used 1.27 grams of hydrogen, but we had 5.00
  grams.
• Therefore, we had 3.73 grams of hydrogen in
  excess.

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Theoretical, actual and percent yields

• Theoretical yield
• The amount of product that should be formed
  according to the chemical reaction.
• Actual yield
• The amount of product actually formed.
• Percent yield
• Ratio of actual to theoretical yield, as a .

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Yield

• Less product is almost always produced than
  expected.
• Possible reasons
• A reactant may be impure.
• Some product is lost mechanically since the
  product must be handled to be measured.
• The reactants may undergo unexpected reactions -
  side reactions.
• No reaction truly has a 100 yield.

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Percent yield

• The amount of product actually formed divided by
  the amount of product calculated to be formed,
  times 100.
• yield
  x 100
• In order to determine yield, you must be able
  to recover and measure all of the product in a
  pure form.

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Yield example

• Example. The final step in the production of
  aspirin is the reaction of salicylic acid with
  acetic anhydride.
• 48.6 g of aspirin is produced when 50.0 g of
  salicylic acid and an excess of acetic anhydride
  are reacted. What is the yield?

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Structural Formulae
C A T A L Y S T