Stoichiometry

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CHAPTER 12

• Stoichiometry

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The Arithmetric of Equations

• What is Stoichiometry?
• Comes from the Greek words stoicheion, meaning
  element, and metron, meaning measure.
• The branch of chemistry that deals with the mass
  relationships of elements in compounds and the
  mass relationships between reactants and products
  in a chemical reaction.

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The Arithmetric of Equations

• 2 types of Stoichiometry
• 1. Composition
• Deals with the mass relationships of elements in
  compounds.
• 2. Reaction
• Involves the mass relationships between reactants
  and products in a chemical reaction.

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The Arithmetric of Equations

• In this chapter we will be discussing Reaction
  Stoichiometry, which is based on chemical
  reactions and the law of conservation of matter.
• All Reaction Stoichiometry calculations start
  with a balanced chemical equation (Number of
  moles).
• A balanced chemical equation provides the same
  kind of quantitative information that a recipe
  does.
• Chemists use balanced chemical equations as a
  basis to calculate how much reactant is needed or
  product is formed in a reaction.

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The Arithmetric of Equations

• A balanced chemical equation can be interpreted
  in terms of different quantities, including
  number of atoms, molecules, or moles mass and
  volume.
• Example
• 2H2(g) O2(g) ? 2H2O(g)
• According to this balanced chemical equation you
  need
• 2 moles of hydrogen and 1 mole of oxygen to
  produce 2 moles of water.
• 4 atoms of hydrogen and 2 atoms of oxygen to
  produce 4 atoms of hydrogen and 2 atoms of
  oxygen.
• 2 molecules of hydrogen and 1 molecule of oxygen
  to produce 2 molecules of water.
• 4 g of hydrogen and 32 g of oxygen to produce 36
  grams of water.

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The Arithmetric of Equations

• If we assume standard temperature and pressure
  (STP 0C and 1 atmosphere), the equation also
  tells you about the volume of gases.
• 1 mole of any gas at STP occupies a volume of
  22.4 Liters.
• So, in the previous equation it tells us that we
  need
• 44.8 L of hydrogen and 22.4 L of oxygen to
  produce 44.8 L of water.

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The Arithmetric of Equations

• Mass and atoms are conserved in every chemical
  equation (Law of Conservation of Mass).
• However, molecules, formula units, moles, and
  volumes are not necessarily conserved.

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The Arithmetric of Equations

• Examples
• Balance the following equations and give the mole
  ratio, mass ratio, and volume ratio between the
  reactants and the products.
• ___N2(g) ___H2(g) ? ___NH3(g)
• ___CO(g) ___H2(g) ? ___CH3OH(l)
• ___CH4(g) ___O2(g) ? ___CO2(g) ___H2O(g)
• ___H2S(g) ___O2(g) ? ___SO2(g) ___H2O(g)
• ___C2H2(g) ___O2(g) ? ___CO2(g) ___H2O(g)
• ___K(s) ___H2O(g) ? ___KOH(aq) ___H2(g)
• ___C2H5OH(l) ___O2(g) ? ___CO2(g) ___H2O(g)

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Chemical Calculations

• Reaction Stoichiometry problems can be classified
  according to the information given in the problem
  and the information you are expected to find, the
  unknown.
• Stoichiometirc problems are solved by using mole
  ratios from the balanced equation to convert the
  given quantity.
• In chemical calculations, mole ratios are used to
  convert between moles of reactants and moles of
  product, between moles of reactants, or between
  moles of products.

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Chemical Calculations

• In a typical stoichiometric problem, the quantity
  given is first converted to moles.
• Then the mole ratio from the balanced chemical
  equation is used to calculate the number of moles
  of the wanted substance.
• Finally, the moles are converted to any other
  unit of measurement related to the unit mole, as
  the problem requires.

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Chemical Calculations
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Chemical Calculations

• Example
• In a spacecraft, the carbon dioxide exhaled by
  astronauts can be removed by its reaction with
  lithium hydroxide, LiOH. How many moles of LiOH
  are required to react with 20 mol of CO2, the
  average amount exhaled by a person each day?
• ___CO2(g) ___LiOH(s) ? ___Li2CO3(s) ___H2O(l)

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Chemical Calculations

• Example
• In photosynthesis, plants use energy from the sun
  to produce glucose, C6H12O6, and oxygen from the
  reaction of carbon dioxide and water. What mass,
  in grams, of glucose is produced when 3.00 mol of
  water react with carbon dioxide?

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Chemical Calculations

• Example
• The first step in the industrial manufacture of
  nitric acid is the catalytic oxidation of
  ammonia. The reaction is run using 824 grams of
  NH3 and excess oxygen. How many moles of NO are
  formed? How many moles of H2O are formed?
• ___NH3(g) ___O2(g) ? ___NO(g) ___H2O(g)

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Chemical Calculations

• Example
• Tin (II) Fluoride, SnF2, is used in some
  toothpastes. It is made by the reaction of tin
  with hydrogen fluoride. How many grams of SnF2
  are produced from the reaction of 30.00 grams of
  HF with Sn?

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Chemical Calculations

• Example
• How many molecules of oxygen are produced when
  29.2 grams of water is decomposed by electrolysis?

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Chemical Calculations

• Example
• Nitrogen monoxide and oxygen gas combine to form
  the brown gas nitrogen dioxide, which contributes
  to photochemical smog. How many liters of
  nitrogen dioxide are produced when 34 L of oxygen
  reacts with an excess amount of nitrogen
  monoxide? Assume conditions of STP.

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Chemical Calculations

• Example
• Phosphorus and hydrogen can be combined to formed
  phosphine (PH3). How many liters of phosphine are
  formed when 0.42 L of hydrogen reacts with
  phosphorus?

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Limiting Reagent and Percent Yield

• In a laboratory a reaction is rarely carried out
  with exactly the required amounts of each of the
  reactants.
• In most cases, one or more reactants is present
  in excess (leftovers).
• Once one of the reactants is used up, no more
  product can be formed.
• The substance that is completely used up first in
  a reaction is called the Limiting Reagent.
• Limiting Reagent The reactant that limits the
  amounts of the other reactants that can combine -
  and the amount of the product that can form - in
  a chemical reaction.
• The reactant that is not completely used up in a
  reaction is called the Excess Reagent.

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Limiting Reagent and Percent Yield

• Example
• Suppose you mix 5 moles of Carbon with 10 moles
  of Oxygen and allow the following reaction to
  take place.
• C(s) O2(g) ? CO2(g)
• What is the mole to mole ratio in this reaction?
• Which reactant is limiting the reaction?
• Which reactant is in excess?

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Limiting Reagent and Percent Yield

• There are 2 steps to finding limiting reactants
• Choose one of the reactants (A) and calculate the
  amount of moles of the other reactant (B) that is
  required.
• Compare the calculated amount with the amount of
  (B) actually available.
• If more is required than available, (B) is the
  limiting reactant.
• If less is required, the reactant that you
  started with (A) is the limiting reactant.

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Limiting Reagent and Percent Yield

• Example
• Silicon dioxide (quartz) is usually quite
  unreactive but reacts readily with hydrogen
  fluoride according to the following equation
• ___SiO2 ___HF ? ___SiF4 ___H2O
• What is the mole ratio in this reaction?
• If 2.0 moles of HF is combined with 4.5 moles of
  SiO2, which is the limiting reactant?

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Limiting Reagent and Percent Yield

• Example
• Some rocket engines use a mixture of hydrazine
  (N2H4) and hydrogen peroxide (H2O2) as propellant
  system according to the following equation
• ___N2H4 ___H2O2 ? ___N2 ___H2O
• What is the mole ratio in this equation?
• Which is the limiting reactant in this reaction
  when 0.75 moles of N2H4 reacts with 0.5 moles of
  H2O2?
• How much excess reactant is left over?
• How much of each product is formed?

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Limiting Reagent and Percent Yield

• The amounts of products calculated in the
  stoichiometric problems in this chapter so far
  represent Theoretical Yields.
• Theoretical Yield The maximum amount of product
  that can be produced from a given amount of
  reactant.
• In most chemical reactions, the amount of product
  obtained is less than the theoretical yield.
• The measured amount of a product obtained from a
  reaction is called the Actual Yield.

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Limiting Reagent and Percent Yield

• The actual yield is what you would actually get
  by doing the experiment in the laboratory.
• Why is the actual yield always less than the
  theoretical yield?
• Some of the reactant might take part in competing
  side reactions.
• Once the product is formed it is usually
  collected in impure form and during the
  purification process some of the product is lost.

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Limiting Reagent and Percent Yield

• The comparison between theoretical and actual
  yields is called the Percent Yield.
• Percent Yield The ratio of the actual yield to
  the theoretical yield, multiplied by 100.
• Percent Yield Actual/Theoretical 100

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Limiting Reagent and Percent Yield

• Example
• Chlorobenzene (C6H5Cl) is used in the production
  of many important chemicals such as aspirin,
  dyes, and disinfectants. One industrial method of
  preparing C6H5Cl , used in the chemical industry,
  is the following reaction
• C6H6 Cl2 ? C6H5Cl HCl
• When 36.8 grams of C6H6 react with an excess of
  Cl2, the actual yield of C6H5Cl is 38.8 grams.
  What is the percent yield of C6H5Cl?

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Limiting Reagent and Percent Yield

• Example
• Methanol can be produced through the reaction of
  CO and H2 in the presence of a catalyst
• ___CO ___H2 ?catalyst? ___CH3OH
• If 75.0 g of CO reacts to produce 68.4 g of
  CH3OH, what is the percent yield of CH3OH?