CP Chemistry Unit 6 Stoichiometry-Test Plan
Key
Learning
1
2
CP Standard/Outcomes
Importance
Stoichiometry
1
Mass Percent, Empirical and Molecular Formula
2
Essential Vocabulary:
 Mass Percent
 Empirical Formula
 Molecular Formula
 Molar ratio
 Limiting reagent
 Excess reactant
 Experimental yield
 Theoretical yield
 Percent yield
Equations, Calculations, and Applications:
 The procedure for finding the empirical and molecular formulas of a compound
 Using mass percent to find the mass of element in a given mass of compound
 Determining the empirical and molecular formulas of a compound from mass %
and molar mass of elements
 Calculate amounts of reactant and product in a balanced chemical equation
 Solution stoichiometry
 Calculate limiting reagent
 Calculate reactant in excess
 Calculate percent yield
exp erimental
% yield 
 100
theoretical
2
Using your textbook as a resource for Unit 6
Topic
Mole ratios and
stoichiometry
Limiting reagents
Percent yield
Reading
Section 3.4 p. 106-108
Section 3.4 p. 110-113
Section 3.4 p. 114-116
3
Skill Practice:
Interpreting a Balanced Chemical Equation
Example:
H2
+
1 molecule
Cl2
→
1 molecule
2 HCl
2 molecules
Complete the blanks:
N2
+
→
3 H2
____ molecule(s)
2 NH3
____ molecule(s)
____ molecule(s)
It follows that any increase of these coefficients will be in the same ratio!
Example:
2 H2
+
__4___ molecule(s)
→
O2
__2__ molecule(s)
2 H2O
__4__ molecule(s)
Complete the blanks:
H2
1.
+
__3__ molecule(s)
N2
2.
3.
3 H2
_20__ molecule(s)
→
+
+
4. __6.02 x 1023__ molecule(s)
→
O2
_1000_ molecule(s)
Cl2
2 HCl
_____ molecule(s)
2 NH3
____ molecule(s)
____ molecule(s)
H2
→
_____ molecule(s)
+
2 H2
Cl2
→
_____ molecule(s)
____ molecule(s)
2 H2O
____ molecule(s)
2 HCl
_____ molecule(s)
4
Since Avogadro’s number = 6.02×1023 molecules = 1 mole
2 H2
5.
+
__2__ mole
→
O2
____ mole
2 H2O
____ mole
The coefficients in a chemical equation give the mole relationships of reactants and
products in a reaction.
Give the mole relationships for each of the following:
H2
6.
____ mole(s)
C3H8
7.
Cl2 →
+
2 HCl
____ mole(s)
+
____ mole(s)
5 O2
→
____ mole(s)
____ mole(s)
3 CO2
____ mole(s)
+
4 H2O
____ mole(s)
Mole Ratios:
2 H2
+
O2
→
2 H2O
A mole ratio is the mole relationship between two specific chemicals:
Example:
 The mole ratio between H2/H2O is: 2 mole H2 = 2 mole H2O
 The mole ratio between H2/O2 is: 2 mole H2 = 1 mole O2
 The mole ratio between O2/H2O is: 1 mole O2 = 2 mole H2O
8. Using the reaction in problem 6 above:
 What is the mole ratio for H2/Cl2? ____________ = ____________
 What is the mole ratio for H2/HCl? ____________ = ____________
 What is the mole ratio for Cl2/HCl? ____________ = ____________
BE SURE TO INCLUDE THE UNITS AND CHEMICAL WHEN WRITING A
MOLE RATIO. NOT JUST NUMBERS!!!!
5
Skill Practice: Using Mole-Mole Relationships
Example:
C3H8
+
5 O2
→
3 CO2 +
4 H2 O
1. How many moles of O2 are needed to completely react with 1 mole of C3H8? Answer:
5 moles
2. How many moles of CO2 form when 5 moles of O2 react? Answer: 3 moles
3. How many moles of H2O form when 1 mole of C3H8 react? Answer: 4 moles
4. How many moles of C3H8 are required to produce 8 moles of H2O?
Answer: 2 moles.
Practice:
2 Mg
+
O2 
2 MgO
1. How many moles of O2 are needed to completely react with 2 moles of Mg? Answer:
________
2. How many moles of MgO form when 3 moles of O2 react? Answer: _____
3. How many moles of MgO form when 6 moles of Mg react? Answer: ______
4. How many moles of O2 are required to produce 20 moles of MgO?
Answer: ________
________________________________________________ Instructor signature
6
Mole Ratio Worksheet
Name: _____________
*Balance all equations first!
1) Given this equation: ___N2 + ___H2 ---> ___NH3, write the following molar
ratios:
a) N2 / H2 _____________ = ______________
b) N2 / NH3 _____________ = ______________
c) H2 / NH3 _____________ = ______________
2) Given the following equation: ___H2 + ___S8 ---> ___H2S, write the following
molar ratios:
a) H2 / H2S _____________ = ______________
b) H2 / S8 _____________ = ______________
c) H2S / S8 _____________ = ______________
3) Answer the following questions for this equation:
___ H2 + ___O2 ---> ___ H2O
a) What is the H2 / H2O molar ratio? _____________ = ______________
b) Suppose you had 20 moles of H2 on hand, how many
moles of H2O could you make? ______________
c) What is the O2 / H2O molar ratio? _____________ = ______________
d) Suppose you had 20 moles of O2, how many moles of H2O
could you make? ____________
4) Use this equation: ___N2 + ___H2 ---> ___NH3 for the following
problems.
a) If you used 1 mole of N2, how many moles of NH3 could be
produced? __________
b) If 10 moles of NH3 were produced, how many moles of N2 would
be required? _________
c) If 3.00 moles of H2 were used, how many moles of NH3 would be
made? ___________
d) If 0.600 moles of NH3 were produced, how many moles of H2 are
required? _________
7
Balanced Equations and Mole Ratios
Name: ____________
1. Balance each equation.
2. Write a mole ratio between the two chemicals with a star (*). Follow the
example in the first problem.
Mole ratio
1.
*
*
__1___NH4NO2 --> __2___H2O
2.
_____H2
3.
_____MgCO3 --> _____MgO + _____CO2
4.
_____P4 + _____Cl2 --> _____PCl5
5.
_____CrO3
6.
_____IF5
7.
_____NH3 + _____O2 --> _____NO + _____H2O
8.
*
*
_____HBrO3 --> _____Br2O5
9.
*
_____NO2 + _____H2O
*
+ ___1__N2
*
+ _____N2 --> _____NH3
*
*
*
*
*
*
--> _____Cr2O3 + _____O2
*
+ _____H2O
*
--> _____HF
+ _____HIO3
*
10.
1 mole NH4NO2 = 2 mole H2O
*
+ _____H2O
*
--> _____HNO3
*
*
_____NH4NO3 --> _____N2O + _____H2O
+ _____NO
8
9
10
11
12
13
14
15
Limiting Reagents 2
Name: _____________________
1. Given the following reaction:
C3H8
(hint: balance the equation first)
+ O2
------->
CO2
+
H2O
If you start with 14.8 g of C3H8 and 3.44 g of O2,
a) determine the limiting reagent
b) determine the number of g of carbon dioxide produced
c) determine the number of grams of H2O produced
d) determine the number of grams of excess reagent left.
2. Given the following equation:
Al2(SO3)3 + 6 NaOH ------> 3 Na2SO3
+ 2 Al(OH)3
a) If 10.0 g of Al2(SO3)3 is reacted with 10.0 g of NaOH, determine the limiting reagent
b) Determine the number of g of Al(OH)3 produced
16
c) Determine the number of grams of Na2SO3 produced
d) Determine the number of grams of excess reagent left over in the reaction
3. Given the following equation:
Al2O3
(hint: balance the equation first)
+ Fe ------> Fe3O4
+
Al
a) If 25.4 g of Al2O3 is reacted with 10.2 g of Fe, determine the limiting reagent
b) Determine the number of g of Al produced
c) Determine the number of grams of Fe3O4 produced
d) Determine the number of grams of excess reagent left over in the reaction
17
Solutions Worksheet
2.
Calculate the mass of water produced when 0.333 L of 0.500 M NaOH is added to 3.0
grams of acetic acid. The equation is:
NaOH(aq) + HC2H3O2(aq)  NaC2H3O2(aq) + H2O(l)
3.
Calculate the number of grams of AgCl formed when 0.200 L of 0.200 M AgNO3 is added
to 75.0 grams of CaCl2. The equation is:
2 AgNO3(aq) + CaCl2(aq)  2 AgCl(s) + Ca(NO3)2(aq)
4.
Calculate the mass of AgCl formed when 0.250 L of 0.100 M solution of NaCl is added to
0.100 L of 0.200 M AgNO3.
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
5.
Calculate the mass of BaSO4 formed when 0.875 L of 0.200 M Na2SO4 solution is added to
0.500 L of 0.500 M BaCl2 solution.
BaCl2(aq) + Na2SO4(aq)  2 NaCl(aq) + BaSO4(s)
18
19
20
21
22
23
24
25
Information: Empirical Formulas
Molecules can be represented by using either a molecular formula or an empirical formula. The
molecular formula tells you exactly how many atoms of each element are in the compound. For
example, in the table below, compound #2 has exactly 4 carbons and 8 hydrogens in each
molecule. Observe the table below that shows four organic molecules along with a molecular and
empirical formula for each one:
Molecule
#1
#2
#3
#4
Molecular Formula
C2H4
C4H8
C3H8
C8H18
Empirical Formula
CH2
CH2
C3H8
C4H9
Critical Thinking Questions
1. What is an empirical formula?
2. How can molecules #1 and #2 have the same empirical formula even though they are different
molecules?
3. Given the empirical formula for a compound is it possible to determine the molecular formula?
If so, explain how.
4. Given the molecular formula for a compound is it possible to determine its empirical formula?
If so, explain how.
5. Give the empirical formula for each of the molecules below:
a) N2O6
b) C2H4O2
c) C4H14
d) C3H5
Information: Percent Composition
Sometimes it is needful to know the composition of a compound. For example, 39.3% of the mass
of sodium chloride is due to sodium. The other 60.7% of the mass is from chlorine. So, in a 100 g
sample of sodium chloride, there are 39.3 g of sodium and 60.7 g of chlorine. This type of data is
known as percent composition. The percent composition tells you the percentage by mass of an
element in a compound. There is a convenient formula for finding the percent composition of an
element in a compound:
26
(obtained from periodic table)
percent composition of element " X" 
mass of x in one mole of the compound
 100
mass of one mole of the compound
(obtained from periodic table)
Let us look at how the percent composition of calcium (Ca) in calcium chloride (CaCl2) was
determined.
percent composition of Ca 
mass of Ca in one mole of CaCl 2
 100
mass of one mole of CaCl 2
from periodic table for calcium
40.1 g
percent composition of Ca 
 100  36.1%
111.1 g
from periodic table for calcium + 2 chlorines;
40.1 + 2(35.5) = 111.1
As another example, consider calculating the percent composition of nitrogen in Ca3N2:
From periodic table for 2 nitrogen atoms: 2(14.0)=28.0
percent composition of N 
28.0 g
 100  18.9%
148.3 g
from periodic table for 3 calcium + 2 nitrogen
atoms: 3(40.1) + 2(14.0) = 148.3
Critical Thinking Questions
Note: For the following questions use 12.0 g/mol for the molar mass of carbon and 1.01 g/mol for
the molar mass of hydrogen. These values can be found on the periodic table.
1. Verify that in C4H10 the percent composition of carbon is approximately 82.6%.
2. Calculate the percent composition of sodium in Na2S.
27
Information: Formulas and Percent Composition
Table 1: Percent composition and formulas of some compounds.
Name
Hexene
Structural Formula
H2C
Propene
CH
CH2
H2C
CH2
CH
CH2
CH3
Molecular
Formula
% Comp.
of H
% Comp.
of C
C6H12
14.4
85.6
CH3
CH
Benzene
HC
CH
HC
CH
C6H6
CH
HC
CH
HC
CH
CH2
CH2
Cyclobutadiene
1,5-hexadi-yne
7.8
HC
C
C
CH
Critical Thinking Questions
3. Verify that the percent composition of C and H given for hexene in Table 1 are correct.
4. Fill in the blanks in Table 1 by determining the percent composition and the molecular
formulas of each compound.
5. Can you determine a compound’s structural formula if you are given the molecular
formula? Explain.
6. What is true about the percent composition of two different compounds that each have the
same ratio of carbon to hydrogen?
7. Can you determine a molecule’s molecular formula solely from the percent composition?
Explain.
28
8. It is possible to complete the following table using only the information in Table 1
without the aid of a calculator or periodic table. Try it! (Hint: consider question #6.)
Molecular Formula
C 8 H8
C10H20
% Composition of H
% Composition of C
An empirical formula is a formula that describes the lowest whole-number ratio of elements in a
compound. An example of an empirical formula is CH, which is the empirical formula for
benzene whose molecular formula was given in Table 1.
Crical Thinking Questions
9. What is the empirical formula of a compound whose percent composition is 92.2% carbon
and 7.8% hydrogen? (See question 8 and Table 1)
10. Verify that the empirical formula for hexene (see Table 1) is CH2.
Information: Calculating the Empirical Formula
When you know the percent composition of each element in a compound, you can calculate the
empirical formula of that compound. The following example will illustrate how to do this.
Example 1: A certain compound is 30.4% nitrogen and 69.6% oxygen by mass. What is
the empirical formula of the compound?
Step #1: Divide each percentage by the molar mass from the periodic table:
For Nitrogen: 30.4  2.17
14.0
For Oxygen: 69.6  4.35
16.0
From the periodic table for nitrogen and oxygen
Step #2: Find the ratio of nitrogen to oxygen. To do this, find the smallest answer
obtained in step #1. In this example, the smallest answer is 2.17. Now divide each
of your answers to step #1 by this smallest number. In this example, you should
divide each answer by 2.17:
For Nitrogen:
2.17
 1.00
2.17
For Oxygen:
4.35
 2.00
2.17
Step #3: Write the formula. The answers from step #2 are the subscripts in the
formula! The formula is NO2.
If in step #2 you get something like Nitrogen = 1.00 and Oxygen = 2.50 then the formula you write
in step #3 would be NO2. 5. This doesn’t make sense because all subscripts must be whole
numbers. You would need to double each subscript. The formula would be N1x2O2.5x2 = N2O5.
11. Find the empirical formula for a compound that contains 82.4% nitrogen and 17.6%
hydrogen.
29
Information: Calculating the Molecular Formula from the Empirical
Formula
Remember that the empirical formula is just a simplification of the molecular formula. For
example, consider the empirical formula NO. There are several possible molecular formulas
including: N2O2, N3O3, N4O4, etc. Which one is it? Notice that the possible formula N2O2 is made
up of two of the empirical formulas, NO. Similarly, N3O3 is made up of three of the empirical
formulas, NO. How do we know which empirical formula is correct? All you need is the molar
mass or molecular mass of the molecular formula.
Critical Thinking Questions
12. The empirical formula for a certain compound is NO. The molar mass of the compound is
60.0 g/mol.
a) What is the molar mass of the empirical formula? (Use the periodic table.)
b) Divide the molar mass of the compound (given in the question) by the molar mass of
the empirical formula found in part a.
c) Your answer to part b tells you how many empirical formulas are in the molecular
formula. You now should be able to write the correct molecular formula, which is
N2O2. Verify that the correct molecular formula is N2O2.
13. a) What is the empirical formula of a compound whose percent composition by mass is
85.7% carbon and 14.3% hydrogen?
b)
If the compound has a molar mass of 56 g/mol, what is the molecular formula?
(Follow the steps from question 12abc.)
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31
Chemistry Can Make Us Rich (Lab)
Name:
_____________
Goals:
 To turn copper into silver!
 To practice using the “mole hill” to calculate how much silver we should get.
 To determine the correct chemical equation for the reaction.
Safety: AgNO3 is toxic and will stain your hands. Wear gloves while making the solution and be
very careful with it.
Pre-Lab: This reaction is a single displacement reaction. AgNO3 is reacting with Cu wire. You
will need to predict the products. Draw the complete reaction twice. Once showing the reaction if
we assume Cu is +1 charge, and a second time if it is +2. Balance both reactions.
Reaction 1:
Reaction 2:
Part 1:
 Make 40 mL of a 0.07M AgNO3. Put the final solution in the smallest beaker that you
have.
 Take (approximately) 8 inches of Cu wire. Coil it over a pencil. Weigh the Cu: ______g
 Hang the Cu wire from a pencil so that most of it is in the AgNO3 solution. Return to your
seat so that we can work on some things while letting the reaction work its magic.
Part 2: Procedure to get your silver clean, dry, and ready to weigh:







Reminder: Careful with the AgNO3!!!
Decant (pour carefully) the solution into one of the waste containers. Be careful to not let
any of the Ag fall out.
Use a water spray bottle to rinse the Cu wire. Do this so that the Ag metal falls with the
water into the beaker.
Decant the liquid into the waste bin being careful not to lose any Ag.
Repeat this process again.
Dip your Cu wire in acetone, let it air dry for 3 minutes, then weigh it on the same balance
that you used yesterday. ______g.
Place your Ag beaker on the tray for it to be dried overnight.
32
Now for the calculations:

Not all of the Cu wire completed the reaction. How can we determine the mass of Cu that
actually did the reaction? (Look at the only two values that you have measure thus far).
Show your work here and put your answer in appropriate spot in the box below:
Work:
g of Ag expected (calculated)
g of Cu
consumed by the
reaction. (this is
measured on
the balance)
______ g Ag using a 1:1 ratio
Actual g of Ag
recovered. (this is
measured on the
balance)
_______ g
________ g
______ g Ag using a 1:2 ratio
Show your calculation here:

Now use the “mole hill” to convert the grams of Cu to grams of Ag. Remember: You
will have to do two separate calculations; one using the mole ratio from reaction 1 and
another using the mole ratio in reaction 2. Put your answers in the box above.

Weigh your dried Ag. Which reaction was better at predicting the amount of Ag
recovered? Circle the correct reaction on the front page.
33
Limiting Reagent Lab
Name:
_____________
Goal: To identify the limiting reactant in a precipitation reaction.
Materials:
 0.29 CaCl2
 solid Na2CO3 (choose an amount between 0.1 and 1.0 grams)
Safety:



Wear goggles until your lab bin is put away.
CaCl2 and Na2CO3 are moderately toxic if ingested and can cause rash on your skin. Be
careful with these reagents.
Wash your hands before leaving class.
Pre-Lab:
1. Predict the products and balance the equation for the reaction that we will be doing
today:
CaCl2
+
Na2CO3

2. What type of reaction are we doing? (Classify the reaction)
_____________________________
3. We are planning to isolate the product of this reaction today. Remember, the
product in these reactions is always the precipitate. Which of the two products is
going to precipitate? ________________(Use your solubility chart.) Record this
answer in the blanks in table 1 and add state symbols to your balanced equation.
Procedure: There is none. Science is about exploring, thinking, trying, messing up, trying again,
and succeeding. You have the knowledge, ingenuity, and background to figure this out on your
own. Remember: Our goal is to react the two reactants to make a precipitate. We will need to
isolate the precipitate to get a mass of the product.
Procedure: (Record what you decided to do here. Bullet points are fine, but include enough detail
so that somebody else could repeat what you did on their own.)
34
Data: This table will guide you through figuring out which reagent was limiting in your
experiment and how much product you should anticipate. Input your measured, and calculated
amounts in the table below. Show your calculations on the bottom of this sheet.
Table 1.
g of
CaCl2
added
to the
beaker
g of
Na2CO3
added
to the
beaker
moles
of
CaCl2
added
to the
beaker
moles
of
Na2CO3
added
to the
beaker
Which
reagent was
limiting?
___________
moles of the
limiting
reagent
moles of
__________
precipitate
expected
g of
__________
precipitate
expected
(calculated)
Actual g
of
________
precipitate
recovered.
(measured)
0.29 g
Calculations:
Post-Lab Problem: Determine how much of the excess reagent was left over after the reaction.
35
Introduction to Stoichiometry Lab
Name: ____________________
Goal: To make 0.5 g of CuO. (And to make an educated guess as to the amount of starting
material that will give us 0.5 g of CuO.)
Safety: Tie back long hair and loose clothing near the Bunsen burner. Copper carbonate is toxic.
Do not ingest it and wear goggles at all times. Be careful with the flame. Turn it off if you are
leaving your station for any amount of time.
Waste: After they cool down, the products can go in the garbage.
Carbonates are a special class of compounds that do this reaction when heated:
CuCO3  CuO + CO2
What kind of reaction is this? ________________ (choose from synthesis,
decomposition, double and single displacement.)
What you need to do: We want to make exactly 0.5 g of the CuO product. To do this we need to
decide how much CuCO3 to start with.

Weigh between 0.5 and 1.5 g of CuCO3 into your evaporating dish.
Mass of evaporating dish empty ________ g
Mass of evaporating dish filled with CuCO3 ________ g
Mass of CuCO3 that you are starting with ________ g

Set up a Bunsen burner as shown below:

Put the evaporating dish on the wire guaze and start heating. Stir the powder with
your glass stirring rod until all of the green has turned to black.
36




Let the dish cool until you can safely touch it. Weigh the evaporating dish again.
________ g
We need to be sure that all of the stuff has changed to product. Heat it again for
another minute, then cool, and weigh the dish again. ________ g
Did the mass change from the last time you weighed it? If so repeat heating until it
doesn’t change by more than 0.1 g. Record the final mass of the evaporating dish
here: _________ g
Mass of empty evaporating dish __________g (You already weighed this above.)
Mass of black CuO after you subtract out the dish: _________ g
Did you get exactly 0.5 g? (0.45-0.55g is okay.) If you didn’t, predict how much CuCO3 you
should start with to get exactly 0.5 g of CuO in the end. Try it again with your new predicted
amount.
Record the data that you collected for your second attempt here:
Mass of CuCO3 at start
Mass of the CuO at the end.
Remember units!
Conclusion:

What mass of CuCO3 do you need to start with to get 0.5 g of CuO at the end? _____g
(SHOW THIS TO YOUR INSTRUCTOR FOR CONFIRMATION.)

Why does the mass change going from reactants to products?