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? 2 Eggs
How much butter is needed for the amount of chocolate chips
used? 1 Cup
How many eggs would we need to make 9 dozen cookies? 6 Eggs
Cookies and Chemistry…Huh!?!?




Just like chocolate chip cookies
have recipes, chemists have
recipes as well
Instead of calling them recipes,
we call them chemical
equations
Furthermore, instead of using
cups and teaspoons, we use
moles to measure amounts
Lastly, instead of eggs, butter,
sugar, etc… we use chemical
compounds as ingredients
Stoichiometry
The Math of Chemical
Equations
Just what is stoichiometry?

The word stoichiometry is
derived from two Greek words:
stoicheion (meaning element)
and metron (meaning measure).
Jeremias Benjamin Richter (1762-1807)
was the first to lay down the principles of
stoichiometry.
In 1792 he wrote: “Stoichiometry is the science
of measuring the quantitative proportions or
mass ratios in which chemical elements stand to
one another.”
This was verified by Antoine Lavoisier
with the Law of Conservation of Mass
When do we use stoichiometry?
Whenever chemists want to know
how much of a desired product will
be produced.
 Chemists always start with a
balanced equation.

How does it work?


The chemist uses the balanced equation
to guide them in calculating the
maximum amount of products.
Since stoichiometry is used to predict
things in a “perfect world”, we can expect
that the reaction in the lab won’t create
the maximum amount of product that
was calculated.
How do I use the chemical
equation?



Looking at a chemical equation tells us
how much of something you need to react
with something else to get a product (like
the cookie recipe)
The chemical equation is written like a
mathematical expression.
When balanced, the coefficients tell how
many moles of each reactant or product is
needed or expected.
The important thing is that the
equation must be BALANCED!!!!
Example:


2H2 + O2 2H2O
One way to understand the information
conveyed by a chemical equation is to
convert the equation into an English
sentence.
2H2(g) + O2(g) 2H2O(l)
2 “parts” hydrogen gas reacts with one “part” oxygen
gas to form 2 “parts” liquid water.
Now, we can get a bunch of relationships – the
“parts” could be: molecules, atoms, moles, or
masses
(the coefficients tell how many moles, atoms or
molecules of each chemical are needed in the
“recipe”)
The most important relationship for stoichiometry is
the MOLE relationship!
What is a mole ratio?
Mole Ratios:
 The mole ratio is based on the balanced
chemical equation.
 You will use the coefficients to get the
mole ratio between two DIFFERENT
substances
 These ratios will be used as conversion
factors for stoichiometry problems.
How do you use a mole ratio?



The mole ratio is an equivalent which means
that you can arrange the ratio in any way
needed.
The mole ratio is the KEY to calculations
based upon a chemical equation.
With it, you can calculate the amount of any
other reactant in the equation and the
maximum amount of product you can obtain.
Molar Ratios




The expressing the ratio of moles in an
equation using the coefficients
Ex. 2H2 + O2
 2H2O
What is the mole ratio between H2 and
O2 ?
2 mol
H2ratios1can
mol
2
What other
mole
we O
come
up
or
with?
1 mol O2
2 mol H2
Warning:



The coefficients of a reaction only give
the ratio in which substances react and
form.
They DO NOT directly tell you, HOW
MUCH MASS is reacting.
You will need to convert from moles to
mass using molar mass as the
conversion factor.
Example of a mole ratio
problem:
N2 + 3H2  2NH3

Write the mole ratios for N2 and H2
1 mol N2
3 mol H2
or
3 mol H2
1 mol N2

Write the mole ratios for NH3 and
2 mol NH3
3 mol H2
H2
or
3 mol H2
2 mol NH3
Stoichiometry has 5 basic steps:
Step 1: Write and balance the
equation
Step 2: Write down all information
that you’re given (Make sure to
identify what you are trying to
find!)
Step 3: Convert everything given into
moles. (Use molar mass as the
conversion factor if you have to.)
Step 4: Use a mole ratio to solve
for what you are trying to find.
(Use the coefficients from the
balanced chemical equation to
get the correct mole ratio.)
Step 5: Convert, if needed,
everything into correct units. (IF
you need to get from moles to
mass use the molar mass of the
substance as the conversion
factor.)
Mole-to-Mole Stoichiometry



One way to change ion ore, Fe2O3, into metallic
iron is to heat it together with hydrogen:
Fe2O3 + 3H2  2Fe + 3H2O
How many moles of iron are made from 25.0
moles of Fe2O3 reacting with excess hydrogen?
Mole-to-Mole
K: 25 moles Fe2O3
 UK: ? moles Fe
 Use Dimensional Analysis :
25 moles Fe2O3  2 mol Fe 

 = 50. mol Fe
 1 mol Fe2O3 

More examples of mole-to-mole
stoichiometry problems:
2H2 + O2  2H2O
1. How many moles of H2O are produced
when 5.00 moles of oxygen are burned in
excess hydrogen gas?
10.0 mol H2O
Examples of a mole-mole
stoichiometric problem
continued:
2H2 + O2  2H2O
2. If 3.00 moles of H2O are produced, how
many moles of oxygen must be consumed
in a reaction with excess hydrogen gas?
1.50 mol O2
Examples of a mole-mole
stoichiometric problem
continued:
2H2 + O2  2H2O
3. How many moles of hydrogen gas must
be used, when reacting with 0.25 mol
oxygen gas to create water.
0.50 mol H2
Mole-to-Mass Stoichiometry


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 (as a conversion factor) 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   70.90 g Cl2 



2
mol
Na

  1 mol Cl2 
= 177g Cl2
Practice

Calculate the mass, in grams, of iodine
required to react completely with 0.50
moles of aluminum to make aluminum
iodide.
2 Al + 3 I2  2 AlI3
Mass-to-Mole Stoichiometry


We can also start with mass and 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
EX: 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 = 0.185
18.0 g H2O 6 mol H20 mol C2H6
Practice


Calculate how many moles of oxygen are
required to make 10.0 g of aluminum oxide
4 Al + 3 O2  2 Al2O3
Mass-to-Mass Stoichiometry

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!)
1. We must go from grams to moles by using
molar mass of the compound given
2. Go from moles of compound given to moles
of compound needed using the mole ratio
from the coefficients
3. Go back to grams of compound we are
interested in by using molar mass of the
compound you are looking for
Mass-to-Mass Stoichiometry
Problems

Ex: Calculate how many grams of
ammonia are produced when you react
2.00g of nitrogen with excess hydrogen.
N2 + 3 H2  2 NH3
Practice

How many grams of calcium nitride are
produced when 2.00 g of calcium reacts
with an excess of nitrogen?
3 Ca + N2  Ca3N2
Gas Stoichiometry (Volume-toVolume) @ STP
1.
2.
3.
4.
Figure out volume of gas given.
Convert to moles using the fact that
22.4 L=1 mole @ STP
Change from moles of compound given to
moles of new compound using molar ratio
from the coefficients.
Change moles of new compound to volume
using 22.4 L = 1 mole
Gases and Stoichiometry
2 H2O2 (g)  2 H2O (g) + O2 (g)
3.5 L of H2O2 was decomposed in a
flask. What is the volume of O2 made
at STP?
Bombardier beetle uses
decomposition of
hydrogen peroxide to
defend itself.
What if we are not just looking for a
volume to volume calculation?


See what the question is asking us and use
appropriate conversion factors.
For example, going back to:
2 H2O2 (g)  2 H2O (g) + O2 (g)
What if you wanted to know how many
moles of water were formed from the
decomposition of 0.25 L of H2O2 at STP?
Another Example

Find the number of grams of oxygen
necessary for the combustion of 6.7 L of
magnesium, assuming the reaction
occurs at STP.
2 Mg + O2  2 MgO
Answer: 4.8 g O2
LIMITING
REACTANT
Limiting Reactant: 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
If we had the specified amount of 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



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
s.
This is called the limiting reactant.
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 Reactant:
Example


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, in
grams, is produced?
2 Al + 3 Cl2  2 AlCl3
Start with Al:
10.0 g Al
1 mol Al
27.0 g Al

Now Cl2:
35.0g Cl2
1 mol Cl2
71.0 g Cl2
2 mol AlCl3 133.5 g AlCl3
2 mol Al
1 mol AlCl3
2 mol AlCl3 133.5 g AlCl3
3 mol Cl2
= 49.4g AlCl3
1 mol AlCl3
= 43.9g AlCl3
LR Example Continued

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 Reactant Practice

15.0 g of potassium reacts with 15.0 g of
iodine. Calculate which reactant is
limiting and how much product is made.
2K + I2  2KI

From the previous question, how much of the
excess reactant is left over?
Percent Yield
Theoretical Yield: The maximum quantity of product that
can be obtained, according to the reaction
stoichiometry, from a given quantity of reactant
Actual Yield: The quantity of product actually obtained
from experimentation
•Percentage Yield = Actual Yield
Theoretical yield
X
100%
What if I get an answer over
100%?



Percent yield can NEVER be over 100%
This would conflict with the Law of
Conservation of Matter
If you obtain a yield over 100% you
either:
• A.
•
Did your calculation incorrectly by putting
theoretical yield over percent yield OR
B. Made a mistake in the experiment
Percent yield problems:


We calculated that 19.6g of KI forms
from 15.0g K and 15.0g of I2.
If we ran the experiment and only
obtained 18.5g of KI what is the percent
yield?
Limiting Reactant & Percent
Yield: Recap
1.
2.
3.
4.
5.
6.
You can recognize a limiting reactant problem
because there is MORE THAN ONE GIVEN
AMOUNT.
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 percent yield, divide the mass of what you
get from the experiment from mass of the product
from the limiting reactant and multiply by 100%.