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STOICHIOMETRY BASICS. Unit 12 Chemistry Langley. **Corresponds to Chapter 12 in the Prentice Hall Chemistry Textbook. INTRODUCTION. Stoichiometry Dervived from two Greek words meaning measured and element 4 types Mole-Mole Mass-Mole Mass-Mass Volume-Volume Mathematically based

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stoichiometry basics


Unit 12



**Corresponds to Chapter 12 in the Prentice Hall Chemistry Textbook

  • Stoichiometry
    • Dervived from two Greek words meaning measured and element
    • 4 types
      • Mole-Mole
      • Mass-Mole
      • Mass-Mass
      • Volume-Volume
    • Mathematically based
    • Balanced equations very important; first place to start
interpreting chemical reaction
Interpreting Chemical Reaction
  • From a balanced chemical equation, you can describe a chemical reaction in terms of quantities of products and reactants. These quantities include the number of molecules, atoms, moles, mass, and volume of the compounds used as well as each element used to make the respective compounds
interpreting a chemical reaction
Interpreting a Chemical Reaction
  • N2 + 3H2 2NH3 (Figure 12.3, pg. 357)
    • Number of atoms: 2 atoms of nitrogen, 6 atoms of hydrogen, 8 atoms of ammonia
    • Number of molecules: 1 molecule of nitrogen, 3 molecules, 2 molecules of ammonia (1:3:2 ratio)
    • Number of relative moles: 1 mole of nitrogen, 3 relative moles of hydrogen, 2 relative moles of ammonia
      • Molecules and moles are not the same thing
    • Mass: 28.0 g of nitrogen, 6.0 g of hydrogen, 34 g of ammonium
    • Volume: 22.4L of nitrogen, 67.2L of hydrogen, 44.8L of ammonia
      • Because 1 mol of any gas at STP occupies a volume of 22.4L
molar mass of compounds
Molar Mass of Compounds
  • The molar mass (MM) of a compound is determined the same way, except now you add up all the atomic masses for the molecule (or compound)
    • Ex. Molar mass of CaCl2
    • Avg. Atomic mass of Calcium = 40.08g
    • Avg. Atomic mass of Chlorine = 35.45g
    • Molar Mass of calcium chloride = 40.08 g/mol Ca + (2 X 35.45) g/mol Cl 110.98 g/mol CaCl2





molar mass calculation
Molar Mass Calculation
  • Calculate the Molar Mass of calcium phosphate
    • Formula = (NH4)3(PO4)
    • Masses elements:
    • Molar Mass =

Atoms or Molecules

Divide by 6.02 X 1023

Multiply by 6.02 X 1023


Multiply by atomic/molar mass from periodic table

Divide by atomic/molar mass from periodic table

Mass (grams)


molar mass Avogadro’s numberGrams Moles particles

Everything must go through Moles!!!

chocolate chip cookies
Chocolate Chip Cookies!!

1 cup butter

1/2 cup white sugar

1 cup packed brown sugar

1 teaspoon vanilla extract

2 eggs

2 1/2 cups all-purpose flour

1 teaspoon baking soda

1 teaspoon salt

2 cups semisweet chocolate chips

Makes 3 dozen

How many eggs are needed to make 3 dozen cookies?

How much butter is needed for the amount of chocolate chips used?

How many eggs would we need to make 9 dozen cookies?

How much brown sugar would I need if I had 1 ½ cups white sugar?

recipes and equations
Recipes and Equations
  • Just like chocolate chip cookies have recipes, chemists have recipes as well
  • Instead of calling them recipes, we call them reaction equations
  • Furthermore, instead of using cups and teaspoons, we use moles
  • Lastly, instead of eggs, butter, sugar, etc. we use chemical compounds as ingredients
chemistry recipes
Chemistry Recipes
  • Looking at a reaction tells us how much of something you need to react with something else to get a product (like the cookie recipe)
  • Be sure you have a balanced reaction before you start!
      • Example: 2 Na + Cl2  2 NaCl
      • This reaction tells us that by mixing 2 moles of sodium with 1 mole of chlorine we will get 2 moles of sodium chloride
      • What if we wanted 4 moles of NaCl? 10 moles? 50 moles?
starting point balanced equation
Starting Point-Balanced Equation
  • Write the balanced reaction for hydrogen gas reacting with oxygen gas.

2 H2 + O2 2 H2O

    • How many moles of reactants are needed?
    • What if we wanted 4 moles of water?
    • What if we had 3 moles of oxygen, how much hydrogen would we need to react and how much water would we get?
    • What if we had 50 moles of hydrogen, how much oxygen would we need and how much water produced?
mole mole calculations
Mole-Mole Calculations
  • These mole ratios can be used to calculate the moles of one chemical from the given amount of a different chemical
  • Example: How many moles of chlorine is needed to react with 5 moles of sodium (without any sodium left over)?

2 Na + Cl2 2 NaCl

5 moles Na 1 mol Cl2

2 mol Na

= 2.5 moles Cl2

mole mole calculations1
Mole-Mole Calculations
  • How many moles of ammonia are produced when 0.60 moles of nitrogen reacts with hydrogen?
    • N2(g) + 3H2 (g)  2NH3(g)

2.0 mol NH3

0.6 mol N2

= 1.2 mol NH3

1.0 mol N2

mass mole calculations
Mass-Mole Calculations
  • Sometimes you are going to start with mass and will have to convert to moles of product or another reactant
  • We use molar mass and the mole ratio to get to moles of the compound of interest
    • Calculate the number of moles of ethane (C2H6) needed to produce 10.0 g of water
    • 2 C2H6 + 7 O2 4 CO2 + 6 H20

10.0 g H2O 1 mol H2O 2 mol C2H6

18.0 g H2O 6 mol H20

= 0.185 mol C2H6

mass mole calculations1
Mass-Mole Calculations
  • Calculate how many moles of oxygen are required to make 10.0 g of aluminum oxide
mole mass calculations
Mole-Mass Calculations
  • Most of the time in chemistry, the amounts are given in grams instead of moles
  • We still go through moles and use the mole ratio, but now we also use molar mass to get to grams
      • Example: How many grams of chlorine are required to react completely with 5.00 moles of sodium to produce sodium chloride?

2 Na + Cl2 2 NaCl

5.00 moles Na 1 mol Cl2 70.90g Cl2

2 mol Na 1 mol Cl2

= 177g Cl2

mole mass calculations1
Mole-Mass Calculations
  • Calculate the mass in grams of Iodine required to react completely with 0.50 moles of aluminum.
mass mass calculations
Mass-Mass Calculations
  • Most often we are given a starting mass and want to find out the mass of a product we will get (called theoretical yield) or how much of another reactant we need to completely react with it (no leftover ingredients!)
  • Now we must go from grams to moles, mole ratio, and back to grams of compound we are interested in
mass mass calculations1
Mass-Mass Calculations
  • Ex. Calculate how many grams of ammonia are produced when you react 2.00g of nitrogen with excess hydrogen.
  • N2 + 3 H2  2 NH3

2.00g N2 1 mol N2 2 mol NH3 17.06g NH3

28.02g N2 1 mol N2 1 mol NH3

= 2.4 g NH3

mass mass calculations2
Mass-Mass Calculations
  • How many grams of calcium nitride are produced when 2.00 g of calcium reacts with an excess of nitrogen?
volume volume calculations
Volume-Volume Calculations
  • 1 mole of any substance = 22.4L
  • 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 of nitrogen monoxide? Assume conditions of STP.
volume volume calculations1
Volume-Volume Calculations
  • 2NO + O2(g)  2NO2(g)
    • Know: volume of oxygen = 34L O2

2 mol NO2/1 mol O2

1 mol O2 = 22.4 L O2 (at STP)

1 mole NO2 = 22.4 L NO (at STP)

34 L O2

1 mol O2

2 mol NO2

22.4 L NO2

= 64 L O2

22.4 L O2

1 mol O2

1 mol NO2

limiting reactants
Limiting Reactants

1 cup butter

1/2 cup white sugar

1 cup packed brown sugar

1 teaspoon vanilla extract

2 eggs

2 1/2 cups all-purpose flour

1 teaspoon baking soda

1 teaspoon salt

2 cups semisweet chocolate chips

Makes 3 dozen

If we had the specified amount of all ingredients listed, could we make 4 dozen cookies?

What if we had 6 eggs and twice as much of everything else, could we make 9 dozen cookies?

What if we only had one egg, could we make 3 dozen cookies?

limiting reactant
Limiting Reactant
  • Most of the time in chemistry we have more of one reactant than we need to completely use up other reactant.
  • That reactant is said to be in excess(there is too much).
  • The other reactant limits how much product we get. Once it runs out, the reaction. This is called the limiting reactant.
limiting reactant1
Limiting Reactant
  • To find the correct answer, we have to try all of the reactants. We have to calculate how much of a product we can get from each of the reactants to determine which reactant is the limiting one.
  • The lower amount of a product is the correct answer.
  • The reactant that makes the least amount of product is the limiting reactant. Once you determine the limiting reactant, you should ALWAYS start with it!
  • Be sure to pick a product! You can’t compare to see which is greater and which is lower unless the product is the same!
limiting reactant2


Limiting Reactant:
  • 10.0g of aluminum reacts with 35.0 grams of chlorine gas to produce aluminum chloride. Which reactant is limiting, which is in excess, and how much product is produced?

2 Al + 3 Cl2 2 AlCl3

  • Start with Al:
  • Now Cl2:

10.0 g Al 1 mol Al 2 mol AlCl3 133.5 g AlCl3

27.0 g Al 2 mol Al 1 mol AlCl3

= 49.4g AlCl3

35.0g Cl2 1 mol Cl2 2 mol AlCl3 133.5 g AlCl3

71.0 g Cl2 3 mol Cl2 1 mol AlCl3

= 43.9g AlCl3

limiting reactant3
Limiting Reactant
  • We get 49.4g of aluminum chloride from the given amount of aluminum, but only 43.9g of aluminum chloride from the given amount of chlorine. Therefore, chlorine is the limiting reactant. Once the 35.0g of chlorine is used up, the reaction comes to a complete .
limiting reactant4
Limiting Reactant
  • 15.0 g of potassium reacts with 15.0 g of iodine. Calculate which reactant is limiting and how much product is made.
finding the amount of excess
Finding the Amount of Excess
  • By calculating the amount of the excess reactant needed to completely react with the limiting reactant, we can subtract that amount from the given amount to find the amount of excess.
  • Can we find the amount of excess potassium in the previous problem?
finding excess
Finding Excess
  • 15.0 g of potassium reacts with 15.0 g of iodine. 2 K + I2 2 KI
  • We found that Iodine is the limiting reactant, and 19.6 g of potassium iodide are produced.

15.0 g I2 1 mol I2 2 mol K 39.1 g K

254 g I2 1 mol I2 1 mol K

= 4.62 g K USED!

15.0 g K – 4.62 g K = 10.38 g K EXCESS

Given amount of excess reactant

Amount of excess reactant actually used

Note that we started with the limiting reactant! Once you determine the LR, you should only start with it!

limiting reactant recap
Limiting Reactant: Recap
  • Convert ALL of the reactants to the SAME product (pick any product you choose.)
  • The lowest answer is the correct answer.
  • The reactant that gave you the lowest answer is the LIMITING REACTANT.
  • The other reactant(s) are in EXCESS.
  • To find the amount of excess, subtract the amount used from the given amount.
  • You can recognize a limiting reactant problem because there is MORE THAN ONE GIVEN AMOUNT.
  • If you have to find more than one product, be sure to start with the limiting reactant. You don’t have to determine which is the LR over and over again!
empirical vs molecular formulas
Empirical vs. Molecular Formulas
  • Empirical Formula of a compound shows the smalles whole number ration of the atoms in the compound
    • A compound is analyzed and found to contain 25.9% nitrogen and 74.1% oxygen. What is the emprical formula of the compound?
      • 25.9 g N * (1 mol/14.0g) = 1.85 mol N
      • 74.1 g O * (1 mol O/16.0g) = 4.63 mol O
      • Empirical Formula = N2O5
empirical vs molecular formulas1
Empirical vs. Molecular Formulas
  • Molecular Formula tells you the actual number of each kind of atom present in a molecule of the compound
  • Calculate the molecular formula of a compound whose molar mass is 60.0 g and empirical formula is CH4N.
    • Find empirical formula mass: 12+4+14=30g
    • Molar mass/efm = 60/30 = 2
    • Molecular Formula = C2H8N2
percent theoretical and actual yield
Percent, Theoretical, and Actual Yield
  • Actual Yield – the amount produced in a reaction (amount of products)
  • Theoretical Yield – the amount that should have been produced in a reaction
    • Maximum amount of product that could be formed from given amount of reactants
  • Percent Yield – (actual/theoretical) X 100%
    • Often less than the theoretical