1. Stoichiometry

2. Defined
Stoichiometry –The study of quantitative (measurable) relationships that exist in chemical formulas and chemical reactions.
You must have:
1. The correct formulas – “criss-cross” or prefixes
2. Balanced Equations – Five types, verify law or conservation of mass

3. Analysis of a chemical reaction/equation
Consider:
N2H4 + 2H2O2 → N2 + 4H2O
Same as saying…
1mole N2H4 + 2moles H2O2 → 1mole N2 + 4 moles H2O
The mole ratio is the “recipe” for the reaction.

4. Mole-Mole Problems
Factor-Label Method – problem solving method based upon treating units in calculations as if they were algebraic factors.
N2H4 + 2H2O2 → N2 + 4H2O
Ex.1.4 moles of N2H4 gives how many moles of N2?
1.4 moles N2H4 x mole N2 = moles N2
mole N2H4

5. Another Example 3Zn + 2H3PO4 → Zn3(PO4)2 + 3H2
How many moles of Zn3(PO4)2 will be produced from 2.18 moles of H3PO4?
2.18 moles H3PO4 x moles Zn3(PO4)2
mole H3PO4
1.09 moles Zn3(PO4)2

6. Verifying the Law of Conservation of Matter
2H2 + O2 → 2H2O
Calculate the mass of the reactants and the mass of the products
2 mol H2 x g H mol O2 x g O2 = 36.04g
1 mol H mol O2
Or 2 mol H2O x g H2O = g
1 mol H2O

7. Mass – Mass Problems Steps Find molar mass of both substances
change mass to moles of given
use molar bridge (moles of given moles of unknown)
change moles to mass of unknown
***Remember, coefficients in the balanced equation represent the relative number of moles of reactants and products!***

8. Mass-Mass Problems
What mass of water is produced from 1.5 grams of glucose?
C6H12O6 + 6 O2 → 6 CO2 + 6 H2O
1st. change mass to moles using molar mass of glucose. (180 g/mol)
2nd. Use molar bridge to change from moles of given to moles of unknown.
3rd change to grams of water

9. The example of Mass-Mass
C6H12O6 + 6 O2 → 6 CO2 + 6 H2O
1.5g C6H12O6 to g of water.
0.9g H2O

10. Another example
Ex. What mass of aluminum oxide is produced when 2.3 g of aluminum reacts with iron (III) oxide?
4.3g Al2O3

11. Mass-Volume problems 1st mass of given 2nd moles of given (molar mass)
3rd moles of unknown (molar ratio)
4th volume of gaseous unknown
If I have 125 g of Al2O3 how many L of Al do I STP?

12. Another Example
Find the mass of aluminum required to produce 1.32L of H2 STP
2Al + 3 H2SO4 → Al2(SO4)3 + 3 H2

13. Volume-volume problems
Same as mole-mole just using volumes instead!
If I have 15.5 L of N2 gas, how many L of H2 will react?
N2 + 3 H2 → 2 NH3

14. 11-3 Limiting Reactants and Percent Yield
When chemicals combine, they are usually in nonstoichiometric proportions. That means there will be a limiting reactant.
Limiting Reactant will be completely used up in the reaction.
The other, leftover amount is said to be in excess.
**The quantities of products formed in a reaction are always determined by the quantity of limiting reactant. **

15. LR continued To find the limiting Reactant:
1st solve 2 separate mass-mass problems.
2nd see which number is smaller; this is your limiting reactant.

16. Example of LR
Ex. 3.5 g of Cu is added to 6.0 g silver nitrate. Find the limiting reactant.
(Note: you can calculate the mass of either product, use the easier one to find the molar mass!)
Cu + 2 AgNO3 → Cu(NO3)2 + 2 Ag

17. You can also determine how much of the substance in excess is left over.
1st determine the LR and then use the amount of the LR and use that to find the mass of the other substance.
Ex. 2.0 g of NaHCO3 reacts with 0.5 g of H3C6H5O7 (citric acid) which is the LR and how much of the other will be left over? What volume of CO2 will be produced?

18. example
3 NaHCO3 + H3C6H5O7 → 3 CO2 + 3 H2O + Na3C6H5O7

19. Percent yield Percent yield = actual yield x 100 Expected
Actual = what you got in lab
Expected = what it was supposed to be

20. Example
A piece of copper with a mass of 5.00g is placed in a solution of AgNO3 . The silver metal produced has a mass of 15.2g. What is the percent yield for this rxn?
1st calculate the mass of the expected silver.
Cu + 2AgNO3  2 Ag + Cu(NO3)2

21. Example continued