Chapter 9 Stoichiometry

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Chapter 9Stoichiometry
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Unit Essential Question

• What numerical information can we find from a
  balanced chemical reaction?

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Lesson Essential Question

• How can we tell the amounts of substances
  involved in a reaction?

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Section 1 Calculating Quantities in Reactions

• Coefficients of balanced equations show
  proportions (ratios) of substances in reactions.
• Similar to amounts in a recipe.
• Use this information to adjust amount of
  product(s) made or reactants needed.
• Relative amounts (coefficients) in equation
  moles.
• Example 2 C8H18 25 O2 ? 16 CO2 18 H2O
• 2 moles of C8H18 react with 25 moles of O2 to
  form 16 moles of CO2 and 18 moles of H2O

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Stoichiometry

• Deals with quantities of substances in chemical
  reactions.
• Compare amounts (moles) and also masses.
• Limiting reactants and percent yield
• can be calculated.
• Can be useful for everyday activities such as
  baking.

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Mole Ratio

• Needed in order to perform calculations!
• Allows you to convert between substances.
• Set up using fractions and conversion factors.
• Use dimensional analysis.
• Get ratios from balanced equation.
• Ex 2H2 O2 ? 2H2O
• Mole ratio examples 2mol H2/1mol O2,
• 1mol O2/2mol H2O, or 2molH2/2molH2O

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Steps to Converting Between Amounts of Substances

• Step 1 Balance the chemical equation.
• Step 2 Convert given amount to moles.
• Remember- always start with your given!
• Step 3 Use a mole ratio from the balanced
  equation to convert from the amount of one
  substance to the amount of another substance.
• Use dimensional analysis for all conversions!
• Step 4 Convert out of moles if asked for grams
  or the number of particles.

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Sample Problem 1

• Consider the reaction for the commercial
  preparation of ammonia
• N2 3 H2 ? 2 NH3
• How many moles of hydrogen are needed to prepare
  312 moles of ammonia?
• Notice that mole ratios convert between
  substances and units stay the same (all moles)!
• Moles of nitrogen needed?

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Sample Problem 2

• What mass of NH3 can be made from 1,221 g of H2
  and excess N2?
• N2 3 H2 ? 2 NH3

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Sample Problem 3

• How many grams of C5H8 form from 1.89 x 1024
  molecules C5H12?
• C5H12 (l) ? C5H8 (l) 2 H2 (g)

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Lesson Essential Question

• How can we calculate the amount of products made
  in a reaction?

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Section 2 Limiting Reactants and Percent Yield

• Limiting reactant substances that limit the
  amount of product that can be formed.
• Limiting reactants run out first!
• Excess reactant substances that are not used up
  in the reaction. There are some left over.
• Do not limit the amount of product that can be
  formed.
• Think back to the sandwich activity- what were...
• LR limiting reagent(s)?
• ER excess reagent(s)?

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Theoretical Actual Yield
This is what we have been calculating in practice
problems so far!

• Theoretical the amount that should be able to be
  produced based on the limiting reactant.
• If everything in the reaction went according to
  plan, and all of the reactant(s) reacted, this is
  how much product should be made.
• This is NOT the same as the actual yield- amount
  that is produced based on an experiment
• Error occurs, so actual yield is less than the
  theoretical yield.
• Can never be greater than the theoretical yield.

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Determining the Limiting Reactant

• You will be given amounts of two or more
  reactants.
• Convert all amounts to moles.
• Use mole ratios in the balanced chemical equation
  to determine the number of moles of a product
  that will be produced.
• Convert all reactant amounts into the same
  product!
• Convert the moles of product to grams of product.
• Do this for each reactant amount given and see
  how much product each produces.
• Whichever produces LESS product is your limiting
  reactant- it will be completely consumed first.

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Sample Problem 5

• Identify the limiting reactant and the
  theoretical yield of phosphorus acid, H3PO3, if
  225 g of PCl3 is mixed with 123 g of H2O
• PCl3 3 H2O ? H3PO3 3 HCl

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Sample Problem 6

• Determine the limiting reactant if 14.0g N2 are
  mixed with 9.0g H2. What is the theoretical yield
  in grams of the product?
• N2 3H2 ? 2NH3

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Application

• Cheapest ingredient excess reactant.
• Most expensive limiting reactant.
• Cost effective!
• Examples
• Cider vinegar
• ER oxygen LR ethanol
• Banana Flavoring
• ER acetic acid LR isopentyl alcohol

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Percent Yield

• Actual yield is always less than the theoretical
  yield.
• LR is not always 100 used up.
• Can be due to human errors, not reacting to
  completion, reverse reactions and/or unwanted
  side reactions occurring, and loss during
  purification.
• Efficiency is another name for percent yield.
• Percent Yield
• (actual yield/theoretical yield) x 100 yield
• If your percent yield is greater than 100, there
  could be unreacted or unwanted materials in your
  product.

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Sample Problem 7

• Using your answers from Sample Problem 6,
  calculate the percent yield if 16.1 g NH3 are
  formed in an experiment.
• N2 3 H2 ? 2 NH3

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Sample Problem 8

• How many grams of CH3COOC5H11 should form if 4808
  g are theoretically possible and the percent
  yield for the reaction is 80.5?