1. Chapter
12
(Chapter 9 in your books)
Stoichiometry

2. Section 12.1
The Arithmetic of Equations
OBJECTIVES:
• Explain how balanced equations apply to both chemistry and everyday life.
• Interpret balanced chemical equations in terms of: a) moles, b) representative
particles, c) mass, and d) gas volume (Liters) at STP.
• Identify the quantities that are always conserved in chemical reactions.
Stoichiometry is…
• Greek for “measuring elements”
Pronounced “stoy kee ahm uh tree”
• Defined as: ___________________________________________________________
_______________________________________________________________________
_
• There are 4 ways to interpret a balanced chemical equation
#1. In terms of Particles
• An Element is made of _____________________
• A Molecular compound (made of only nonmetals) is made up of ________________
(Don’t forget the diatomic elements)
• Ionic Compounds (made of a metal and nonmetal parts) are made of _____________
Example: 2H2 + O2 → 2H2O
• Two molecules of hydrogen and one molecule of oxygen form two molecules of
water.
• Another example: 2Al2O3 → 4Al + 3O2
Now read this: 2Na + 2H2O → 2NaOH + H2
#2. In terms of Moles
• The coefficients tell us how many moles of each substance
2Al2O3 → 4Al + 3O2

3. 2Na + 2H2O → 2NaOH + H2
• Remember: A balanced equation is a Molar Ratio
#3. In terms of Mass
• The Law of Conservation of Mass applies
• We can check mass by using moles.
2H2 + O2 → 2H2O
In terms of Mass (for products)
2H2 + O2 → 2H2O
The mass of the reactants must equal the mass of the products.
#4. In terms of Volume
• At STP, 1 mol of any gas = 22.4 L
2H2 + O2 → 2H2O
NOTE: mass and atoms are ALWAYS conserved - however, molecules, formula units,
moles, and volumes will not necessarily be conserved!

4. Practice:
• Show that the following equation follows the Law of Conservation of Mass (show the
atoms balance, and the mass on both sides is equal)
2Al2O3 → 4Al + 3O2

5. Name ___________________________________ Date __________________________
12-1 Section Review
1. Balance this equation: C2H5OH (l) + O2 (g) → CO2 (g) + H2O
a. interpret the equation in terms of numbers of molecules. (Write out a word
equation and include the numbers of molecules of each reactant and product)
b. Interpret the equation in terms of numbers of moles. (Write a word equation and
include the numbers of moles of each reactant and product)
c. Using the molecular weights on your periodic table, show that the balanced
equation obeys the law of conservation of mass (interpret the equation in terms of
masses of reactants and products.

6. 2. Use the equation below to answer the questions that follow.
2K (s) + 2H2O(l) → 2KOH (aq) + H2
a. interpret the following equation in terms of relative numbers of representative
particles
b. interpret the following equation in terms of numbers of moles
c. interpret the following equation in terms of masses of reactants and products

7. Section 12.2
Chemical Calculations
OBJECTIVES:
• Construct “mole ratios” from balanced chemical equations, and apply these ratios in
mole-mole stoichiometric calculations.
• Calculate stoichiometric quantities from balanced chemical equations using units of
moles, mass, representative particles, and volumes of gases at STP.
Mole to Mole conversions
2Al2O3 → 4Al + 3O2
• each time we use 2 moles of Al2O3 we will also make 3 moles of O2
These are the two possible conversion factors to use in the solution of the problem.
Mole to Mole conversions
• How many moles of O2 are produced when 3.34 moles of Al2O3 decompose?
2Al2O3 → 4Al + 3O2
If you know the amount of ANY chemical in the reaction, you can find the amount of
ALL the other chemicals!
Practice:
2C2H2 + 5 O2 → 4CO2 + 2 H2O
• If 3.84 moles of C2H2 are burned, how many moles of O2 are needed?
• How many moles of C2H2 are needed to produce 8.95 mole of H2O?

8. • If 2.47 moles of C2H2 are burned, how many moles of CO2 are formed?
Steps to Calculate Stoichiometric Problems
1. _________________________________________________________________
_
2. _________________________________________________________________
_
3. _________________________________________________________________
_
4. _________________________________________________________________
_
5. _________________________________________________________________
_
Mass-Mass Problem:
• 6.50 grams of aluminum reacts with an excess of oxygen. How many grams of
aluminum oxide are formed?
Another example:
• If 10.1 g of Fe are added to a solution of Copper (II) Sulfate, how many grams of
solid copper would form?

9. Volume-Volume Calculations:
• How many liters of CH4 at STP are required to completely react with 17.5 L of O2 ?
Avogadro told us:
• Equal volumes of gas, at the same temperature and pressure contain the same number
of particles.
• Moles are numbers of ______________________________
• You can treat reactions as if they happen _______________ at a time, as long as you
keep the temperature and pressure the same.
Shortcut for Volume-Volume?
• How many liters of CH4 at STP are required to completely react with 17.5 L of O2?
Note: This only works for Volume-Volume problems.

10. Name ___________________________________________ Date __________________
12-2 Section Review
1. Isopropyl alcohol (C3H7OH) burns in air according to this equation:
2C3H7OH (l) + 9O2 (g) → 6CO2 (g) + 8H2O (g)
a. calculate the moles of oxygen needed to react with 3.40 mol C3H7OH
b. Find the mole of each product formed when 3.40 mol C3H7OH reacts with oxygen
2. What ratio is used to carry out each conversion?
a. mol CH4 to g CH4
b. L CH4 (g) to mol CH4 (g) (at STP)
c. Molecules CH4 to mol CH4
3. The combustion of acetylene gas is represented by this equation?
2C2H2 (g) + 5O2 (g) → 4CO2 (g) + 2H2O (g)
a. How many grams of CO2 and grams of H2O are produced when 52.0 g C2H2 burns

11. b. How many grams of oxygen are required to burn 52.0 g C2H2?
c. Use the answers from a and b to show that this equation obeys the law of
conservation of mass
4. Tin (II) fluoride, formally found in many kinds of toothpaste, is formed in this
reaction:
Sn (s) + 2HF (g) → SnF2 (s) + H2 (g)
a. how many liters of HF are needed to produce 9.40 L H2 at STP
b. How many molecules of H2 are produced by the reaction of tin with 20.0 L HF at
STP
c. How many grams of SnF2 can be made by reacting 7.42 x 1024 molecules of HF
with tin?

12. Section 12.3
Limiting Reagent & Percent Yield
OBJECTIVES:
• Identify the limiting reagent in a reaction.
• Calculate theoretical yield, percent yield, and the amount of excess reagent that
remains unreacted given appropriate information.
• Calculate theoretical yield, percent yield, and the amount of excess reagent that
remains unreacted given appropriate information.
“Limiting” Reagent
• The ____________________________________ is the reactant you run out of first.
• The ____________________________________ is the one you have left over.
• The limiting reagent determines how much product you can make
How do you find out which is limited?
• The chemical that makes the ___________________ amount of product is the
“limiting reagent”.
• You can recognize limiting reagent problems because they will give you
___________________________________________________
• Do two stoichiometry problems, one for each reagent you are given.
Example:
If 10.6 g of copper reacts with 3.83 g sulfur, how many grams of the product (copper (I)
sulfide) will be formed?

13. Another example:
• If 10.3 g of aluminum are reacted with 51.7 g of CuSO4 how much copper (grams)
will be produced?
How much excess reagent will remain?
What is yield?
• Yield is the amount of __________________________________________________.
• There are three types:
1. ___________________________________- what you actually get in the lab when the
chemicals are mixed
2. ___________________________________________- what the balanced equation tells
should be made
3. ___________________________________ =
Example:
• 6.78 g of copper is produced when 3.92 g of Al are reacted with excess copper (II)
sulfate.

14. • What is the actual yield?
• What is the theoretical yield?
• What is the percent yield?
Details on Yield
• Percent yield tells us how “efficient” a reaction is.
• Percent yield can not be bigger than 100 %.
• Theoretical yield will always be larger than actual yield!
Why?

15. Name ______________________________________ Date _______________________
12-3 Section Review
1. What is a limiting reagent?
2. What is an excess reagent?
3. What is the percent yield if 4.65 g of copper is produced when 1.87 g of
aluminum reacts with an excess of copper(II) sulfate?
2Al (s) + 3CuSO4 (aq) → Al2(SO4)3 (aq) + Cu (s)
4. What is the difference between an actual yield and a theoretical yield?
Which yield is larger for a given reaction? How are these values used to
determine percent yield?
5. How many grams of SO3 are produced when 20.0 g FeS2 reacts with 16.0
g O2 according to this balanced equation?
4FeS2 (s) + 15O2 (g) → 2Fe2O3 (s) + 8SO3 (g)