1. Review:
Mole Conversions: Convert 3 mols Oxygen to grams: Convert 42 grams Chlorine to mols: What is % composition? What is the %comp of magnesium in magnesium nitrate?

2. Stoichiometry
Stoichiometry refers to the calculations of chemical quantities from balanced chemical equations.
Where do we use this? Stoichiometry is used to determine the amount of reactant or product necessary to complete a reaction.
Why is it important? This can be used in factories to make sure the appropriate amount of product is made with the least amount of waste.

3. Practical problems
Let’s first look at stoichiometry in our lives. What makes up a grilled cheese? ___Bread slices + ___Cheese ___ Grilled Cheese How many slices of bread are needed to make 24 Grilled cheese sandwiches? How many grilled cheese can be made with 12 slices of cheese?
Butter

4. Interpreting Chemical Equations
The first thing that must be done is to balance the equation!
N2 (g) H2 (g)  NH3 (g)
Here are the kinds of information you can get from the equation:
____ mole N ____ moles H2  ____ moles NH3
___ molecule N ___ molecules H2  ___ molecules NH3
____ liter N ____ liters H2  ____ liters NH3
____ grams N ____ grams H2  ____ grams NH3

5. Mole-Mole Conversions
When we are given an amount of moles in a reaction, we can find the amounts of the other reactants/products.
A mole ratio is needed to convert from one mole to another. We find the mole ratio in the balanced equation.
The mole conversion factor comes from the coefficients of the balanced chemical equation.
Step 1: Write down the what you know.
Step 2: Use a mole ratio to convert.

6. Practice Problems: N2 (g) + 3H2 (g)  2NH3 (g) 1) How many moles of ammonia can be made from 7 moles of nitrogen reacting with an excess of hydrogen? 2) How many moles of hydrogen are required to completely react with 8 moles of nitrogen to produce ammonia? 3) How many moles of hydrogen are needed to react with an excess of nitrogen to make 10 moles of ammonia?

7. What If you are given MASS???
Mass-Mass: (grams  moles  moles  grams)
Step 1: Write down what you know.
Step 2: Use Molar mass to convert to mols.
Step 3: Use a mole ratio from your balanced equation.
Step 4: Use molar mass to convert to grams.
N2 (g) H2 (g)  2NH3 (g)
Practice Problem: How many grams of ammonia can be made from reacting 39.0 grams of nitrogen with an excess of hydrogen?

8. Other Conversion Problems
Volume-Mass: (liters  moles  moles  grams)
Step 1: Write down what you know.
Step 2: Use (1 = 22.4 L) to convert to mols.
Step 3: Use a mole ratio from your balanced equation.
Step 4: Use molar mass to convert to grams.
N2 (g) H2 (g)  2NH3 (g)
Practice Problem: How many grams of ammonia can be made from reacting 20.0 liters of hydrogen with an excess of nitrogen?

9. Other Conversion Problems
Mass-Volume: (grams to moles to moles to liters)
Step 1: Write down what you know.
Step 2: Use molar mass to convert to mols.
Step 3: Use a mole ratio from your balanced equation.
Step 4: Use (1 = 22.4 L) to convert to Liters.
N2 (g) H2 (g)  2NH3 (g)
Practice Problem: How many liters of ammonia can be made from reacting 36.0 grams of nitrogen with an excess of hydrogen?

10. Other Conversion Problems
Volume-Volume: (exactly like mole-mole)
Step 1: Write down the what you know.
Step 2: Use a mole (But with Liters as units) fraction to convert.
N2 (g) H2 (g)  2NH3 (g)
Practice Problem: How many liters of ammonia can be made from reacting 125 liters of hydrogen with an excess of nitrogen?

11. Percent Yield
The Percent Yield is a ratio that tells us how efficient a chemical reaction is.
- The higher the % yield, the more efficient the reaction is. (Unless you go over 100%)
Actual Yield
Theoretical Yield
The actual yield is the amount you experimentally produce when you run the reaction in a lab.
The theoretical yield is the amount you are ideally produce if everything goes perfectly. You can calculate this amount using stoichiometry!
% Yield =
x 100

12. How to Calculate Percent Yield
Step 1: Calculate the theoretical yield for the problem. (This means doing the stoichiometry for the problem)
Step 2: Divide the actual yield by the theoretical yield.
Step 3: Multiply by 100

13. Practice Problem: 2H2 (g) + O2 (g)  2H2O (g)
1) A student reacts 40 grams of hydrogen with an excess of oxygen and produces 300 grams of water. Find the % yield for this reaction.

14. Practice Problem: 2H2 (g) + O2 (g)  2H2O (g)
2) If 2.0 grams of hydrogen completely reacted with 16.0 grams of oxygen but only produced 17.5 grams of water, what is the % yield for the reaction?

15. FOR THE FIRST 5 MINUTES! Get into your lab groups from last week.
Work TOGETHER on the following Problem
a. Write the equation for the reaction of iron (III) phosphate with sodium sulfate to produce iron (III) sulfate and sodium phosphate.
 b. If I perform this reaction with 25 grams of iron (III) phosphate and an excess of sodium sulfate, how many grams of iron (III) sulfate can I make?   
c. If 18.5 grams of iron (III) sulfate are actually made, what is my percent yield?
Sulfate: SO4-2
Phosphate: PO4-3

16. Limiting Reactant
The limiting reactants is the reactant that runs out first. It “limits” the reaction from continuing (when it is used up).
The reactant that is in abundance is called the excess reactant.

17. How to find the Limiting Reactant
Step 1: If given grams, convert reactants to Mols of a product.
Step 2: Compare the amount of product. The reactant that makes less product is the limiting reactant.
If I give you a quantity for both reactants, you must find the limiting reactant before finding the final amount of product.
If I say THE HYDROGEN IS USED COMPLETELY…
What does this mean?

18. Practice Problems:___N2 (g) +___H2 (g)  ___NH3 (g) 1) If 2
Practice Problems:___N2 (g) +___H2 (g)  ___NH3 (g) 1) If 2.7 moles of nitrogen reacts with 6.3 moles of hydrogen, which will you run out of first? 2) If 3.9 moles of nitrogen reacts with 12.1 moles of hydrogen, what is the limiting reagent?

19. How to calculate Excess Reactant
How many moles of excess reagent do you have?
PRE-STEP 1: You must know which reactant is in excess!
Step 1: Start with the amount of product will be produced from the limiting reactant.
Step 2: Convert this to mols of the reactant in excess.
Step 3: Subtract this amount from the original amount.

20. Practice Problem: N2 + 3H2  2NH3
1) If 2.7 moles of nitrogen reacts with 6.3 moles of hydrogen, how much excess reactant will you have?
2) If 3.9 moles of nitrogen reacts with 12.1 moles of hydrogen, how much excess reactant will you have?