Oakland Schools Chemistry Resource Unit
Moles & Stoichiometry Regina Chu South Lyon High School
South Lyon Community Schools 1
Moles & Stoichiometry
Content Statements:
C4.1x:
Compounds have a fixed percent elemental composition. For a compound, the
empirical formula can be calculated from the percent composition or the mass of each
element. To determine the molecular formula from the empirical formula, the molar
mass of the substance must also be known.
C4.6x:
The mole is the standard unit for counting atomic and molecular particles in terms of
common mass units.
C5.2x:
A balance chemical equation will allow one to predict the amount of product formed.
Content Expectations:
C4.1a: Calculate the percent by weight of each element in a compound based on the
compounds formula.
C4.1b: Calculate the empirical formula of a compound based on the percent by weight
of each element in the compound.
C4.1c: Use the empirical formula and molecular weight of a compound to determine
the molecular formula.
C4.6a: Calculate the number of moles of any compound or element given the mass of
the substance.
C4.6b: Calculate the number of particles of any compound or element given the mass
of the substance.
C5.2d: Calculate the mass of a particular compound formed from the masses of
starting materials.
C5.2e: Identify the limiting reagent when given the masses of more than one product.
C5.2g: Calculate the number of atoms present in a given mass of element.
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Instructional Background Information:
C4.1x: Molecular and Empirical Formulae
C4.1a: Calculate the percent by weight of each element in a compound based on the
compound formula.
C4.1b: Calculate the empirical formula of a compound based on the percent by weight
of each element in the compound.
C4.1c: Use the empirical formula and molecular weight of a compound to determine
the molecular formula.
Compounds have a fixed percent composition. To determine the percent by mass (the
percent composition) of each substance in a compound, one must start with the molar
mass of the compound. The mass of each substance is divided by the total mass of the
compound and multiplied by 100 to obtain the percent composition. The benchmark
states that students should be able to determine percent composition for compounds
containing multiple elements and hydrates. Examples of both follow.
% of element =
mass of element
X 100
mass of compound
EX: Determine the percent composition of CuSO4.
The molar mass of the compound is 159.5 grams/mole. The molar mass of copper is
63.5 grams/mol, the molar mass of sulfur is 32.1 grams/mole, and the molar mass of
oxygen is 16 grams/mole. Dividing the molar mass of each individual element by the
total mass of the compound yields the following percentages: percent copper is 39.8%,
percent sulfur is 20.1%, and percent oxygen is 40.1%.
EX: Determine the percent composition of CuSO4•5H20.
The molar mass of CuSO4 is 159.5 grams/mole. The mass of 5 H20 molecules is 90
grams/mole. To determine the percent composition of water in the hydrate, simply
divide the mass of water by the mass of the copper (II) sulfate plus the mass of the
water and then multiply by 100. This gives 36.1%. To determine the percent of each
of the elements individually, simply divide the molar mass of the elements by the total
mass of the compound and multiply by 100.
Empirical and molecular formulae for a compound can be determined from the percent
composition of a compound. It is assumed that you start with a 100 gram sample;
simply remove the percent sign and add a gram sign to find the mass of each
substance that you have. Each substance’s mass is then divided by its molar mass to
determine the number of moles of each substance present. The molar ratio of the
substances can then be obtained by dividing by the smallest number of moles present;
the ratios must be whole numbers, so it might be necessary to multiply the ratios to
obtain whole numbers. If one number is multiplied to get a whole number, all the
ratios must be multiplied. These molar ratios then become the subscripts for the
empirical formula. The benchmark states that students should be able to determine
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percent composition for compounds containing multiple elements and hydrates.
Examples of both follow.
Ex: Determine the empirical formula of a compound if its percent
composition is 24.75% K, 34.77% Mn, and 40.51% O.
24.75 grams K
1mole K
=
0.633 moles K
=
0.633 moles Mn
=
2.53 moles O
39.1 grams K
34.77 grams Mn
1 mole Mn
54.9 grams Mn
40.51 grams O
1 mole O
16.0 grams O
0.633 = 1.000
0.633 = 1.000
2.53 = 4.00
0.633
0.633
0.633
The empirical formula of this compound is KMnO4.
EX: The empirical formula of a hydrate is determined as follows:
A 100.0 gram sample of FeCl3 • ? H2O is heated and 43.72 grams of water are given
off, leaving 56.28 grams of anhydrous salt (FeCl3). To find the number of water
molecules we must first go to moles of each:
43.72 grams H2O
1mole H2O
=
2.426 moles H2O
=
0.3467 moles FeCl3
18.02 grams H2O
56.28 grams FeCl3 1 mole FeCl3
162.35 grams FeCl3
Now divide by the smallest number of moles:
2.426 =
0.3467
6.998 moles of H2O
0.3467 = 1.000 moles FeCl3
0.3467
Therefore, the empirical formula of this hydrate is FeCl3 • 7 H2O.
Using the empirical formula mass and the molecular mass given in the problem, the
molecular formula can be determined. The molecular mass is divided by the empirical
formula mass to obtain a whole number. The subscripts of the empirical formula are
then all multiplied by that whole number to obtain the subscripts for the molecular
formula.
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EX: Find the molecular mass of KMnO4. The empirical formula of the
compound was determined above.
The molecular mass of the compounds is 158 grams. The empirical formula mass is
158 grams. To find the molecular formula, divide the molecular mass by the empirical
formula mass to get a whole number. Multiply all of the subscripts in the empirical
formula by the whole number ratio to determine the molecular formula. In this case,
the whole number ratio is 1, so the empirical formula and the molecular formula are the
same thing.
EX: A compound of N and O has a molecular mass of 95.0 grams. Calculate
the empirical and molecular formula if the compound contains 1.52 g N and
3.47 g O.
1.52 grams N
1mole N
=
0.109 moles N
=
0.217 moles O
14.0 grams N
3.47 grams O
1 mole O
16.0 grams O
0.109 = 1.000
0.217 = 1.991
0.109
0.109
The empirical formula of this compound is NO2 and the molecular formula is N2O4.
C4.6x: Moles
C4.6a: Calculate the number of moles of any compound or element given the mass of
the substance.
C4.6b: Calculate the number of particles of any compound or element given the mass
of the substance.
The mole (named after the unit’s founder—Avogadro) is a standard unit of counting
particles, atoms, and molecules in chemistry, similar to the unit of the dozen. A mole
contains 6.022 x 1023 things—cows, pennies, people, molecules, atoms, etc. Because
the molar masses of different compounds differ so greatly, the only reasonable way to
compare the number of particles of each is by using a like unit, the mole. The molar
mass of any object equals the mass of 6.022 x 1023 pieces of that object. For example,
if one cow weighs 1,000 kg, one mole of cows (6.022 x 1023 cows) weighs 6.022 x 1026
kg.
Why use the mole? One mole of protons or neutrons (which are very similar in mass)
equals approximately 1 gram. The mole is simply a convenient way to mass different
substances.
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EX: Calculate the number of moles found in 23 grams of NaCl.
The molar mass of NaCl can be found on the Periodic Table. Na has a molar mass of
23 grams/mol and Cl has a molar mass of 17 grams/mol; adding those together gives a
molar mass for the compound of 40 grams/mol.
23grams NaCl
1 mol NaCl
=
0.58 moles of NaCl
40 grams
Converting the number of particles or molecules to the number of moles is a similar
conversion; an example follows.
EX: How many molecules are present if you start with 0.84 moles of NaCl?
0.84 mol NaCl
6.022 x 1023 molecules NaCl
=
5.06 x 1023 molecules NaCl
1 mol NaCl
The molar mass is related to the number of particles in that substance through
Avogadro’s Number, 6.022 x 1023.
EX: How many molecules of NaCl do you have if you mass out 32.7 grams?
32.7 grams NaCl
1 mol NaCl
6.022 x 1023 molecules NaCl
40 grams NaCl
1 mole NaCl
4.92 x 1023
=
molecules
C5.2x: Balancing Equations
C5.2d: Calculate the mass of a particular compound formed from the masses of
starting materials.
C5.2e: Identify the limiting reagent when given the masses of more than one product.
C5.2g: Calculate the number of atoms present in a given mass of element.
Coefficients in a balanced equation represent moles of the substances. The ratios
between coefficients (the molar ratios) can be used in calculations. If the amount of
starting material used is known, the amount of product formed can then be calculated.
Once the amount of product formed is known, the limiting and excess reactants can be
determined. The limiting reactant (sometimes called the limiting reagent) is the
substance that is completely consumed during the chemical reaction and thus stops the
reaction. The excess reactant is the reactant that is not completely consumed during
the reaction and thus has some left over after the reaction stops.
EX: The reaction between nitrogen monoxide (NO) and oxygen to from
nitrogen dioxide (NO2) is a key step in photochemical smog formation.
2NO + O2 Æ 2NO2
A) How many moles of NO2 are formed by the complete reaction of 0.254 moles
of O2?
0.254 mol NO2
1 mol O2
=
0.127 mol O2
2 mol NO2
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B) How many grams of NO2 are formed by the reaction of 1.44 grams of NO with
excess O2?
Step 1: Get to moles
Step 2: Set up a mole-to-mole ratio from the balanced equation.
Step 3: Get to units asked for in question
1.44 g NO
1 mol NO
1 mol NO2
46.0 g NO2
30.0 g NO
1 mol NO
1 mol NO2
=
2.21 g NO2
Determining the limiting and excess reactants are as simple as the above examples with
a few extra steps.
EX: The reaction between aluminum and iron (III) oxide generates high
temperatures and is therefore used in the welding process.
2 Al + Fe2O3 Æ 2 Fe + Al2O3
In one process 124 g of Al are reacted with 601 g of Fe2O3. Use the following:
A) Start with one reactant and turn it into the other reactant (either one is fine to
start with). Compare the answer with the given amount in the problem to
determine if you have more of less of what is given. From this information, the
limiting and excess reactants can be determined.
B) Determine how many grams of the product can be formed, starting with the
limiting reactant.
C) Determine how much of your excess reactant is left over.
Part A)
124 g Al
1 mol Al
1 mol Fe2O3
159.7 g Fe2O3
27.0 g Al
2 mol Al
1 mol Fe2O3
=
367 g Fe2O3
The amount of iron (III) oxide needed to react with all of that aluminum is less that
what is stated in the problem (601 g in the problem is more than the 367 g determined
to be needed using the stoichiometric ratios), therefore iron (III) oxide is the excess
reactant and aluminum is the limiting reactant.
Part B)
124 g Al
1 mol Al
1 mol Al2O3
102 g Al2O3
27.0 g Al
2 mol Al
1 mol Fe2O3
=
234 g Al2O3
=
367 g Fe2O3
There are 234 grams of Al2O3 produced during this reaction.
Part C)
124 g Al
1 mol Al
1 mol Fe2O3
159.7 g Fe2O3
27.0 g Al
2 mol Al
1 mol Fe2O3
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Al is the limiting reactant, so we do a stoichiometric conversion to determine how much
Fe2O3 is consumed during the reaction. 367 grams are used, so we simply subtract to
determine how much of the excess reactant is left over.
601g Fe2O3 – 367 g Fe2O3 = 234 g Fe2O3 left over
8
Terms and Concepts
Stoichiometry
Moles
Molar Mass
Empirical Formula
Molecular Formula
Hydrate
Percent Composition
Balanced Equations
Chemical Equations
Limiting Reactant
Excess Reactant
Percent Yield
Theoretical Yield
Avogadro’s Number
Coefficient
Subscript
Mole Ratios
Molarity
Molality
Molecular Weight
Compound
Element
SI units
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Instructional Resources
Sources for background information:
Chang, Raymond. General Chemistry: The Essential Concepts. 5. New
York: McGraw Hill HE, 2008.
Print resources:
HSSCE Companion Document
Suggested online resources for instruction and/or teacher background:
http://www.moleday.org
http://www.chemtutor.com
http://www.hschem.org
http://www.chemistry.org
http://www.boshf.org/chembank/
http://www.nclark.net/
http://misterguch.brinkster.net/chemfiestanew.html
http://dwb4.unl.edu/index.html
http://www.davis.k12.ut.us/staff/ruyahne/firstworksheets.html
http://educ.queensu.ca/~science/main/concept/gen/g09/g09main.htm
http://www.geocities.com/dschan77/mycomputer
http://ed.fnal.gov/arise/guides/chemhot.pdf
http://www.terrificscience.org/index.jsp
http://www.chemtutor.com
http://filebox.vt.edu/c/ckeel/lps/stoich.htm
http://moodle.oakland.k12.mi.us/clarenceville/
www.myteacherpages.com/webpages/TVERESH
www.101science.com
http://filebox.vt.edu/c/ckeel/lps/stoich.htm
http://websites.mhsmi.org/Staff/kakitzmann/
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Moles & Stoichiometry
Activity #1 – Mole and Molar Mass
Questions to be investigated
What is a mole? How are a mole and molar mass related?
Objectives
Students will be able to relate the mass of a given substance to the number of
atoms, molecules, and moles present in the sample. They will also be able to
understand the concept of mole and molar mass.
Teacher Notes
When using this worksheet, the “Mole Machine” worksheet, “The Double Y
Diagram” worksheet, and the “Guide Sheet for Moles” worksheet can also be
passed out for reference for the students; all are shown below.
Materials
Student Worksheet, Periodic Tables, Mole Machine worksheet (optional), Guide
Sheet for Moles (optional), Double Y Diagram (optional)
Real-World Connections
A mole is a term of measurement, much like a dozen is a measurement term for
12 items. Items are frequently packaged and shipped in large amounts: nuts
and bolts at a hardware store are sold by the number or items, but packaged by
the mass. Computers have information stored in bytes and kilobytes; these are
simply measurement tools for the amount of information stored.
Sources
http://www.pogil.org/downloads/Foundations/Mole_Molar_Mass.pdf
(Molar Mass Worksheets)
http://www.chem.vt.edu/RVGS/ACT/notes/Study_Guide-Moles_Problems.html
(Guide Sheet for Moles; tweaked from the website)
http://members.aol.com/rtccpu/DOCs/stoich/DoubleYDiagram.PDF
(Double Y Diagram)
Procedure/Description of Lesson
See the next 4 pages. Clean copies can be printed from the websites. For
optional activities see the 3 pages following the lesson description.
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12
13
14
15
Volume
Volume
In Liters
In Liters
÷ 22.4 L/mol
÷ Molar Mass
Mass
In grams
÷ 6.02x1023
Molecules
Particles
Atoms
X 22.4 L/mol
MOLE
MACHINE
X Molar Mass
Mass
In grams
X 6.02X1023
Molecules
Particles
Atoms
16
17
Guide Sheet for Moles Problems
I. Calculating Molar Mass
1. Multiply the atomic mass of each element by the number of atoms of
that element in the formula (shown by the subscript)
2. Find the sum of all the atomic masses—this is the formula mass (unit is
a.m.u.)
3. Express the formula mass in grams (unit is grams/mol). This is the
Molar Mass.
II. Calculating % Composition (from formula)
1. Find the sum of all the atomic masses (the formula mass).
2. Divide the total atomic mass of each element by the formula mass of the
compound and multiply by 100.
III. Calculating % Composition (from masses of each element)
1. Divide the mass of each element by the total mass of the compound and multiply by 100.
IV. Calculating Empirical Formula (from % Composition)
1. Convert the percent composition of each element to grams based on 100 grams of the
compound. (meaning: take off the % sign and add the units of grams)
2. Convert to moles.
3. Divide each number of moles by the smallest amount of moles.
4. All ratios must be whole numbers. If needed, multiply to obtain whole numbers; all ratios must
be multiplied by the same amount.
V. Calculating Empirical Formula (from experimentally determined masses)
1. Convert to moles.
2. Continue with steps 3 & 4 from IV above.
VI. Finding Molecular Formulas (when molar mass is known)
1.
2.
3.
4.
Calculate the empirical formula.
Use the equation: (empirical formula mass) x = molecular mass
Find the value for x: x = molecular mass/empirical formula mass
Multiply each subscript in empirical formula by the value for x
18
Moles & Stoichiometry
Activity #2 – ChemQuest 31
Questions to be investigated
What is percent composition? How is it determined? Can you use percent composition to
determine the empirical formula of a compound and, if so, how is it determined?
Objectives
Students will be able to write a definition of percent composition and a mathematical
formula for its determination. They will also be able to determine empirical formula from
percent composition.
Teacher Notes
This can be used in place of lecture notes for the assigned topics or on days when you have
a sub. It can also be given as homework the night before the topic is introduced in class.
Materials
Student worksheets, Periodic Tables
Real-World Connections
Percent Composition is similar to test scores for students. Different metal alloys are used
for different things; the percent of each metal in the alloy is essential for creating the right
alloy for the specific job. Jewelry designers use percent composition when using different
carat weights of gold and sterling silver. Pharmacists and pharmaceutical companies use
percent composition when determining the ratio of each compound in medicine.
Construction workers use percent composition when mixing concrete (some mixes are
stronger than others) and when using steel (to determine the % carbon). Different
mineral’s chemical formulas are examples of empirical formulas. News reports/medical
information sheets from pharmacies and pharmaceutical companies mention chemical
formulas (this helps you to know what drugs are made of).
Sources
Neil, Jason. ChemQuest CD. www.ChemistryInquiry.com
CDs of the ChemQuests and ChemActivities (practice worksheets) can be purchased from
the website.
Procedure/Description of Lesson
See the following 4 pages. The lessons can also be printed from the CD, along with the
answer keys.
Assessment Ideas
The CD contains ChemActivities that can be used as assessments or as homework.
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20
21
22
23
Moles & Stoichiometry
Activity #3 – S’mores Lab
Questions to be investigated
What is a limiting reactant? How can you determine your limiting reactant and your excess
reactants?
Objectives
Students will formulate a definition of a limiting reactant and an excess reactant. They will
also be able to determine the limiting and excess reactants in future problems.
Teacher Notes
This lab requires S’mores material for students. It can be run in partners or groups. If a
brand-new, never-been-used-for-chemicals Bunsen burner is used, the S’mores can be
heated and consumed. If you have no brand-new Bunsen burner available, the S’mores
should be eaten cold or warmed in a food-only microwave. This activity can be used as an
introduction to limiting reactants, letting the students determine their own definition for a
limiting reactant, or as a supplement during lecture/notes on the topic.
Materials
Enough Graham crackers, marshmallows, and chocolate for all students; brand-new Bunsen
burner for heating the S’mores; student worksheets
Safety Concerns
S’mores should only be heated over a never-used-for-chemicals Bunsen burner or in a foodonly microwave.
Real-World Connections
Limiting reactants are used when cooking food (hot dogs are sold in packages of 10, while
hot dog buns are sold in packages of 8), making sandwiches, making hamburgers, etc.;
when the Mentos and Diet Coke reaction is too small, more Mentos (the limiting reactant)
should be added. Chemists determine the limiting reactant to maximize the reactions when
using very expensive chemicals (cancer and AIDS drugs).
Sources
Lab Idea taken from:
www.sci.ccny.cuny.edu/~chemwksp/Smores.doc
Procedure/Description of Lesson
See the following 2 pages.
24
Assessment Ideas
A snowball quiz could be given the next day. A regular quiz is given; students answer the
quiz questions on a half-sheet of paper, but do not write their name on it. The class is then
divided in half and told to stand on opposite sides of the room. The quiz is crumpled into a
snowball and a snowball fight ensues. When time is called (usually a few seconds),
students grab one piece of paper from the floor; answers are given and problems worked
out for the students. Students are able to learn from their own mistakes and the mistakes
of the other students that they are grading, but there is no point value associated with this
assessment. Students have a good grasp of where they are with the material and what
they need to work on, but do not have to worry about being penalized if they do not fully
grasp the concept at that moment.
25
Name______________________________
Hour___________
Date_______________
What is a limiting reactant?
A delicious treat known as a S’more is constructed from the following ingredients:
2 graham crackers
1 chocolate bar
1 marshmallow
Suppose we find that these ingredients are available only in full packages, each of which contains
one dozen of the item. The packages of ingredients have the following weights:
graham crackers
chocolate bars
marshmallows
200.0 g
145.0 g
75.0 g
Each group will build S’mores out of the packages of ingredients that you receive. Build as many
S’mores as you can, but don’t eat the S’mores yet.
Questions: (You may use your S’mores to help you visualize these problems)
1. Using G for the graham crackers, C for the chocolate bars, and M for the marshmallows, write
an equation that would represent the production of S’mores from the starting materials.
2. Based on the information given about the three ingredients, which weighs the most? Which
weighs the least? Explain your reasoning.
3. If we have 12 graham crackers (one package), how many chocolate bars and how many
marshmallows do we need to make S’mores with all the graham crackers? Explain.
4. How many S’mores would we be able to make?
5. Suppose we have one package of each of the ingredients. How many S’mores can we make?
Will any of the ingredients be left over? How much?
6. If we make S’mores from the materials described in #5, which ingredient limits the amount of
S’mores you can make—i.e., which ingredient will you run out of first? (This item is known to
chemists as the limiting reactant because it is the reactant that limits the amount of the final
product that can be made)
26
7. How many dozen S’mores will you have made?
8. Suppose you had 500.0 grams of graham crackers, 500.0 grams of chocolate bars, and 500.0
grams of marshmallows.
a. Which item do you have the most of? Which item do you have the least
of? Explain your reasoning.
b. If you attempt to make S’mores with these amounts of ingredients, what
item will you run out of first?
c. How many S’mores can you make?
d. How many grams of each left-over item will you have?
9. Is it correct to say that if we start with 4 lb each of G, C, and M, we should end up with 3 x 4 =
12 lb of S’mores? If not, why not?
Now let’s apply the same concepts to a chemical situation:
Ammonia (NH3) can be formed from the elements N2 and H2, as shown below.
N2
+
3 H2 -------------- >
2 NH3
1. How many moles of ammonia can be made from 1.0 mole of N2 and 3.0 moles of H2?
2. Suppose we had 3.0 moles each of the N2 and H2 available to react. Which of the reactants
would be the limiting reactant?
3. How many moles of ammonia could we make? Would any of the reactants be left over? How
many moles?
4. How many moles of ammonia could we make from 1.0 mole each of N2 and H2?
5. What mass of ammonia could we make from 100.0 grams each of N2 and H2?
27
Moles & Stoichiometry
Activity #4 – Twenty Questions Packet
Questions to be investigated
How are moles, molar mass, and grams related?
Objectives
The student will be able to determine moles of any substance given the starting grams.
The student will be able to convert from grams of reactant to grams of product.
Teacher Notes
The worksheets go in order of difficulty. It is a great resource for drill and practice and a
deeper understanding of the math involved with the mole.
Materials
Student worksheets, Periodic Tables
Real-World Connections
Masses and volumes of products are extremely important in key reactions such as the
deployment of air bags in vehicles, carbon dioxide production to help bread rise, and the
production of ammonia gas for industries.
Sources
Adapted from Kathy Kitzmann, Mercy High School, Farmington Hills, MI
Procedure/Description of Lesson
See the following 3 pages.
Assessment Ideas
These worksheets can be used as an assessment at the end of a unit.
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31
Moles & Stoichiometry
Activity #5 – ChemActivity 31
Questions to be investigated
Can a molecule be identified by its percent composition?
Objectives
Using the percent composition, students will be able to determine the empirical formula of a
compound.
Teacher Notes
This can be used in place of notes on percent composition and empirical formula or as a
supplement to basic notes on these subjects.
Materials
Student worksheets, Periodic Tables
Real-World Connections
Percent Composition is similar to test scores for students. Different metal alloys are used
for different things; the percent of each metal in the alloy is essential for creating the right
alloy for the specific job. Jewelry designers use percent composition when using different
carat weights of gold and sterling silver. Pharmacists and pharmaceutical companies use
percent composition when determining the ratio of each compound in medicine.
Construction workers use percent composition when mixing concrete (some mixes are
stronger than others) and when using steel (to determine the % carbon). Different
mineral’s chemical formulas are examples of empirical formulas. News reports/medical
information sheets from pharmacies and pharmaceutical companies mention chemical
formulas (this helps you to know what drugs are made of). Limiting reactants are used
when cooking food (hot dogs are sold in packages of 10, while hot dog buns are sold in
packages of 8), making sandwiches, making hamburgers, etc.; when the Mentos and Diet
Coke reaction is too small, more Mentos (the limiting reactant) should be added.
Sources
www.pogil.org/downloads/ChemAct_31_4thEd.pdf
Procedure/Description of Lesson
See the following 3 pages.
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33
34
35
Moles & Stoichiometry
Activity #6 – Determination of an Empirical Formula
Questions to be investigated
What is the empirical formula of a compound? How is the empirical formula of a compound
related to its molar mass?
Objectives
The student will determine the empirical formula of an unknown compound from its percent
composition and atomic masses.
Teacher Notes
See the following page titled “Determination of an Empirical Formula Teacher
Notes/Background” for details on the experiment. It is a clever idea that works very well.
Make sure to fill the vials before the students come in. Tell students that they all have
different substances (be careful not to let on that they actually are all the same compound),
but they may appear similar, so they will not get the same results as their neighboring lab
groups. This activity could follow a traditional CuSO4 • 5H2O lab; the copper turns from a
blue to white when all the water is driven off. In this lab, the compound stays white
throughout. Be sure to stress to students that they need to heat the compound for long
enough to drive all of the “gas” off or their results will be skewed.
Materials
Vials numbered 1-12, CaSO4 • 2H2O, goggles, aprons, Bunsen burners, evaporating dish,
balance, tongs, scoop
Safety Concerns
Be careful when using the Bunsen burner or open flame. Remind students that the dish is
extremely hot when heated, so use tongs when removing it from the flame and be careful
not to touch it until cooled.
Real-World Connections
Different mineral’s chemical formulas are examples of empirical formulas. News
reports/medical information sheets from pharmacies and pharmaceutical companies
mention chemical formulas (this helps you to know what drugs are made of).
Sources
Experiment devised by Dave Tanis, based on idea from Jim Seurynck, Hope Chemistry Institute,
July 19, 1995
Procedure/Description of Lesson
See the following 4 pages.
Assessment Ideas
This could be used at or near the end of the unit, so a quiz or a test could follow.
36
Determination of an Empirical Formula
Background:
The law of definite proportions, first proposed by Joseph Proust in 1799, states that a pure
compound always consists of the same elements combined in the same proportions by mass. This
law has been verified by experimental results many times and is generally accepted as correct.
In this experiment, you will be attempting to determine the empirical formula of an
unknown compound AxBy. You will be given a sample of the compound and the molar mass of its
constituents A and B. Heating your compound will cause it to decompose according to the
equation:
AxBy Æ xA(s) + yB(g)
Since “B” is released as a gas, strong heating will leave only “A” of the unknown compound in the
evaporating dish. By comparing the masses before and after heating the relative quantities of A
and B can be determined and by use of their respective atomic masses (molar masses) the relative
number of moles which have combined can be determined. From this ratio, you can calculate the
empirical formula or simplest formula.
There are two types of formulas for a compound. The first type is called the empirical
formula, which is the simplest formula possible for the compound. It consists of the simple, whole
number ratios of the atoms. The molecular formula represents the actual ratio of atoms in a
compound.
Each lab group will receive a different white compound as their unknown material. It is
IMPORTANT to write the unknown number in your data table and record the masses from the
board.
Procedure:
1. Clean a porcelain evaporating dish with soap and water. Dry. Set up a Bunsen burner and
ring stand, put wire gauze on the ring. Set the evaporating dish on the wire gauze and heat
for 3 minutes.
2. Carefully remove the dish with tongs and set on the counter to cool.
3. Once cool, mass the evaporating dish on the balance and record the mass in your data
table.
4. Add approximately 3.0 grams of your unknown to the evaporating dish. Record this mass.
5. Begin heating gently. Increase heat slowly until you have heated strongly for 10 minutes.
This will turn “B” into a gas and drive it off.
6. Remove the dish with tongs and allow to cool. Measure the mass of the dish and contents
(now just “A”) and record.
7. Reheat with a hot flame for about 4 minutes, cool, and mass again. If the mass is not
within 0.01g of the value in step 6, repeat until you get a constant mass.
8. Place the solid residue in the waste container, and then wash the dish, your counter, and
your hands.
37
Empirical Formula of a Compound
Data Table:
Name__________
Date______ Hr___
Unknown #_____
Mass of evaporating dish_________
Mass of dish + unknown _________
Mass of dish + A
_________
Molar mass of A _________
Molar mass of B _________
Observations of unknown before, during, and after heating:
Calculations:
1. Determine the mass of the unknown.
2. Determine the mass of “A”.
3. Determine the moles of “A”.
4. Determine the mass of “B”.
5. Determine the moles of “B”.
6. Determine the empirical formula – mole ratio of A:B
Error Analysis:
Identify 2 sources of error in this lab (remember poor math skills and incorrect balances are not
options)
1.
2.
38
Questions:
1. Calculate the percent composition of your sample.
2. Explain the difference between an empirical and molecular formula.
3. What is the empirical formula of a compound that has
A) 0.36 mol H & 0.090 mol C?
B) 87.5% N and 12.5% H
Discussion/Conclusion:
Discuss the results of this procedure and how scientists can determine the formulas of unknown
compounds. Why do we need to use moles in these calculations?
39
Determination of an Empirical Formula
Teacher Notes/Background
Teacher Background
Each lab group will be given an unknown binary compound to analyze. Actually each group is
given a sample of the same hydrate, CaSO4 • 2H2O. This hydrate loses 1.5 H2O at 128 °C and 2
H2O at 163 °C.
Tell each group that they are receiving an unknown and that the general formula of each group’s
unknown is AxBy. Heating their compound will cause it to decompose according to the following
equation:
AxBy Æ xA(s) + yB(g)
Since “B” is released as a gas, strong heating will leave only the “A” of the unknown in the
crucible.
Each lab group will be told their unknown number. The students should be warned to be sure to
record the unknown number in their report.
Tell each lab group their unknown number. Then assign the gram atomic masses of “A” and “B”
for their unknown using the table below.
Unknown Number
1
2
3
4
5
6
7
8
9
10
11
12
Atomic Mass of “A”
70.4
181.4
113.4
134.0
151.0
159.9
182.0
29.3
105.8
70.3
198.5
116.5
Atomic Mass of “B”
55.9
16.0
(Give this info
45.0
to students)
35.45
79.9
126.9
32.1
31.0
14.0
46.5
21.0
41.1
40
Expected Results (These are the results the students should get.)
Unknown #
Formula
Unknown #
Formula
1
A3B
7
A2B3
2
AB3
8
A4B
3
A3B2
9
AB2
4
AB
10
A5B2
5
A2B
11
A2B5
6
A3B
12
A4B3
Reference
Experiment devised by Dave Tanis, based on idea from Jim Seurynck, Hope Chemistry Institute, July 19, 1995
41
Moles & Stoichiometry
Activity #7 – Limiting Reactant
Questions to be investigated
What is a limiting reactant and how is it determined?
Objectives
Students will be able to write a definition of a limiting reactant and describe how it is
determined. Students will be able to perform calculations to determine the limiting and
excess reactants.
Teacher Notes
This activity is perfect for a conceptual chemistry class. Go to www.pogil.org , click on
Curriculum Materials, then Downloadable Activities, then Collected Activities, then you need
to register for a password. This is a free site filled with numerous guided-inquiry activities.
Materials
Limiting Reactant Worksheets, Periodic Tables
Real-World Connections
Limiting reactants are used when cooking food (hot dogs are sold in packages of 10, while
hot dog buns are sold in packages of 8), making sandwiches, making hamburgers, etc.;
when the Mentos and Diet Coke reaction is too small, more Mentos (the limiting reactant)
should be added. Chemists determine the limiting reactant in a chemical reaction to
maximize the reaction and send it to completion when using very expensive chemicals
(cancer and AIDS drugs).
Sources
www.pogil.org/downloads/Foundations/Limiting_Reactants.pdf
Procedure/Description of Lesson
See the following 6 pages.
42
43
44
45
46
47
48
Moles & Stoichiometry
Activity #8 – Introduction to the Mole
Questions to be investigated
What is a mole? How are mass, molar mass, and the mole related?
Objectives
The students will understand the concept of the mole. The students also will be able to
convert between mass, moles, and molecules.
Teacher Notes
This activity could be used as an opener to the unit. It could be used in place of
lecture/notes for this concept.
Materials
Introduction to the Mole worksheets, Periodic Tables
Real-World Connections
A mole is a term of measurement, much like a dozen is a measurement term for 12 items.
Items are frequently packaged and shipped in large amounts: nuts and bolts at a hardware
store are sold by the number or items, but packaged by the mass. Computers have
information stored in bytes and kilobytes; these are simply measurement tools for the
amount of information stored.
Sources
http://www.terrificscience.org/index.jsp
Procedure/Description of Lesson
See the following 21 pages
49
50
51
52
53
54
55
56
57
58
59
60
61
62
63
64
65
66
67
68
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70
Moles & Stoichiometry
Activity #9 – Micro Mole Rocket Lab
Questions to be investigated
How do stoichiometric ratios (mole ratios) affect chemical reactions? What is a limiting
reactant and how does it affect mole ratios and chemical reactions?
Objectives
The purpose of this experiment is to generate hydrogen and oxygen and determine the
optimum ratio for their combustion reaction to give water. The optimum ratio will be used
to calculate the mole ratio for the reaction of hydrogen and oxygen in a balanced chemical
equation. The concept of limiting reactants will be used to explain the results obtained with
various hydrogen/oxygen gas mixtures.
Teacher Notes
See the notes pages labeled “Micro Mole Rockets—Teacher Guide”
Materials
Hydrochloric acid, HCl, 3 M, 15 mL
Hydrogen peroxide, H202, 3%, 15 ml
Yeast suspension, 2%, 5 ml
Zinc, mossy, Zn, about 5 g
Graduated cylinder, 10-ml
Marker (permanent pen)
One-hole rubber stoppers, to fit test tubes, 2
Test tube rack
Test tubes, small, 2
Pipets, Beral-type, graduated, 1
Scoopula
Safety matches or Piezo sparker
piezoelectric sparking mechanism (modified Charcoal Lighter such as the Aim ‘n Flame
Torch lighter)
electric tape
6-D common nail
film canister
hot glue gun and glue
10 cm of 24 AWG speaker hook-up wire (double solid, not stranded: Radio Shack #2781509)
wire strippers
Safety Concerns
The ignition and launching of the rockets should be closely monitored for student’s safety.
Goggles must be worn throughout the experiment (including the launching). For extra
precaution, a glove can be worn for ignition. All other general lab safety protocols should
be followed as well. Set aside an area for launching the rockets. Set up a target and do
not let students in that area while rockets are being launched.
71
Real-World Connections
A mole is a term of measurement, much like a dozen is a measurement term for 12 items.
Items are frequently packaged and shipped in large amounts: nuts and bolts at a hardware
store are sold by the number or items, but packaged by the mass. Computers have
information stored in bytes and kilobytes; these are simply measurement tools for the
amount of information stored. Limiting reactants are used when cooking food (hot dogs are
sold in packages of 10, while hot dog buns are sold in packages of 8), making sandwiches,
making hamburgers, etc.; when the Mentos and Diet Coke reaction is too small, more
Mentos (the limiting reactant) should be added. Chemists determine the limiting reactant in
a chemical reaction to maximize the reaction and send it to completion when using very
expensive chemicals (cancer and AIDS drugs).
Sources
http://www.csd509j.net/cvhs/kirscha/MicroMoleRockets%200708.doc
http://teacherknowledge.wikispaces.com/Nelson+Inquiry+Lesson
(Teacher info found here)
Additional info found in Flinn ChemTopic Labs “Molar Relationships & Stoichiometry” book,
Vol. 7
Procedure/Description of Lesson
See the following 10 pages.
72
Micro Mole Rockets
Hydrogen and Oxygen Mole Ratio
As adapted from Flinn ChemTopic- Labs - Molar Relationships & Stoichiometry
Introduction
The combustion reaction of hydrogen and oxygen is used to produce the explosive energy needed to
power the space shuttle. The reaction is also being engineered to serve as a source of continuous
energy for fuel cells in electric vehicles. What factors determine the explosiveness of the reaction of
hydrogen with oxygen? In this lab, we will generate microscale quantities of hydrogen and oxygen
and test their explosive nature, first separately, then in mixtures of various proportions. The goal-to
find the most "powerful" gas mixture and use it to launch a rocket across the room!
Concepts
• Mole ratio
• Combustion
Stoichiometry
Limiting reactants
Background
Hydrogen, the most abundant element in the universe, is a colorless, odorless gas. It is combustible,
which means that it burns quite readily. Hydrogen gas is conveniently generated in the lab by the
reaction of zinc metal with hydrochloric acid.
Oxygen, the most abundant element on Earth, is also a colorless, odorless gas. Oxygen gas supports
combustion, that is, it must be present for combustible materials to burn. Small scale quantities of
oxygen gas are conveniently generated in the lab by the decomposition of hydrogen peroxide. The
decomposition reaction of hydrogen peroxide requires a catalyst to initiate the reaction. A variety of
different catalysts, including manganese, manganese dioxide, potassium iodide, and even yeast, have
been used in this reaction. In this lab, yeast will be used to catalyze the decomposition of hydrogen
peroxide and generate oxygen gas.
Experiment Overview
The purpose of this experiment is to generate hydrogen and oxygen and determine the optimum ratio
for their combustion reaction to give water. The optimum ratio will be used to calculate the mole ratio
for the reaction of hydrogen and oxygen in a balanced chemical equation. The concept of limiting
reactants will be used to explain the results obtained with various hydrogen/oxygen gas mixtures.
Pre-Lab Questions
1. Write the balanced chemical equation for the single-replacement reaction of zinc and hydrochloric
acid to generate hydrogen gas. What is the total amount of hydrogen gas that could be produced in
liters with 5.00 grams of zinc? You can assume you have an unlimited amount of hydrochloric
acid?
2. Write the balanced chemical equation for the yeast-catalyzed decomposition of hydrogen peroxide
to generate oxygen gas and water. Note: Since a catalyst is not really a reactant or product, it is
usually written over the arrow. What is the total amount of oxygen gas that will be produced in
liters if you started with 15 grams of hydrogen per oxide?
73
Materials
Hydrochloric acid, HCI, 3 M, 15 mL
Hydrogen peroxide, H202, 3%, 15 ml
Yeast suspension, 2%, 5 ml
Zinc, mossy, Zn, about 5 g
Graduated cylinder, 10-ml
Marker (permanent pen)
One-hole rubber stoppers, to fit test tubes, 2
Test tube rack
Test tubes, small, 2
Pipets, Beral-type, graduated, 1
Scoopula
Safety matches or Piezo sparker
Safety Precautions
Hydrochloric acid is toxic by ingestion and inhalation and is corrosive to skin and eyes. Hydrogen
peroxide is a skin and eye irritant. Avoid contact of all chemicals with skin and eyes and notify your
teacher immediately in the case of a spill. Wear chemical splash goggles and chemical-resistant gloves
and apron. Wash hands thoroughly with soap and water before leaving the laboratory.
Procedure
Construct Gas Generators
1. The gas generators consist of a small test tube, a rubber stopper, a gas delivery tube, and a gas
collection bulb. See Figure 1a.
2. Cut four Beral-type pipets as shown in Figure 1b to obtain four gas-collecting bulbs and four
gas-delivery tubes. Discard the middle part of the pipet stem. It is important that the pipet bulbs have
similar lengths. Trim the lengths so they are equal.
3. Place the gas delivery tube ends into the tops of rubber stoppers as shown in Figure 1a
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4. Prepare a hydrogen gas generator by placing about four pieces of mossy zinc into the bottom of a
small test tube marked “HCl”.
5. Prepare an oxygen gas generator by placing about 2 mL of yeast suspension into the bottom of the
other small test tube marked “H2O2”.
6. Set the test tubes in a test tube rack.
Calibrate Gas Collection Bulbs
7. Fill a 250-ml, beaker about one-half full with tap water.
8. Immerse one of the cut-off pipet bulbs under water. Fill the bulb completely with water and remove it
from the beaker.
9. Squeeze the water out of the pipet bulb into an empty graduated cylinder to measure the total volume
(V) of water in the bulb.
10. Divide the pipet bulb into six, equal-volume increments by following steps 11-12.
11, Refill the pipet bulb, and then squeeze out one-sixth of the total volume (V/6) into an empty
graduated cylinder. Release the squeeze and use a permanent pen to mark the water-level on the side
of the bulb.
12. Squeeze out a second V/6 volume, mark the level again, and repeat for the remainder of the water.
This should serve to divide the bulb into six, equal-volume increments.
13. Once the first pipet bulb has been calibrated, the rest can be copied to save time. Simply rest a wood
splint across the bulb, with the end of the splint flush with the end of the bulb, and mark off the splint
at the same places that the bulb is marked. Then use the splint as a template to mark the rest of the
bulbs.
Collect and Test Hydrogen and Oxygen Gases
14. Add 3 M hydrochloric acid to the mossy zinc in one of the hydrogen gas generators until the liquid
level is about 1 cm below the mouth of the test tube. Cap the tube with the gas delivery stopper.
Note: Wait about one minute before proceeding to step 15. This will allow time for the air to be
purged from the test tube.
15. Completely fill a marked pipet bulb with water and place the bulb over the gas delivery tube to
collect the gas by water displacement. As the bubbles enter the pipet bulb, the water will flow out of
the bulb and down the sides of the test tube to the paper towels.
16. As soon as the bulb is filled with hydrogen, remove it from the gas delivery tube and place a finger
over the mouth of the bulb to prevent the collected gas from leaking out.
17. Hold the gas bulb so the opening is pointed upward and have a classmate quickly strike a match over
the opening of the bulb. After the match is lit, let the hydrogen gas escape into the flame. Record the
results of this "pop-test" in the data table.
75
18. Add 3% hydrogen peroxide to the yeast suspension in one of the oxygen gas generators until the
liquid level is about I cm below the mouth of the test tube. Cap the tube with the gas delivery
stopper. Note: Wait about one minute before proceeding to step 19.
19. Repeat steps 15-17 to collect oxygen gas and test its properties. Record the results of its "pop-test" in
the data table.
Collect and Test Oxygen/Hydrogen Gas Mixtures
20. Completely fill a marked pipet bulb with water and place it over the oxygen gas generator to collect
oxygen.
21. When the bulb is one-sixth full of gas, quickly remove it from the oxygen tube and place it over the
hydrogen gas generator.
22. Continue collecting hydrogen until the bulb is filled with gas. This bulb should contain a 1:5 ratio of
oxygen and hydrogen.
23. Remove the bulb, cap it with a finger, and determine its relative loudness in the "poptest," as
described above for hydrogen and oxygen. Develop a scale to describe how loud this mixture is
compared to pure hydrogen and pure oxygen. Record the result in the data table.
24. Repeat steps 20-23 to collect and test other volume ratios (2:4, 3:3, 4:2, 5:1) of oxygen and hydrogen
(see the data table). Always collect oxygen first, followed by hydrogen. Record all results in the data
table.
25. Rank the gas mixtures on a scale from zero to 10 to describe their relative loudness in the "pop-test."
Let the most "explosive" mixture be a 10, the least reactive gas a zero.
26. Collect various gas mixtures as many times as necessary to determine the optimum ratio of oxygen
and hydrogen for combustion. Note: The pop-test is obviously subjective, but by repeating it several
times with each possible mixture, it should be possible to determine the most explosive (loudest) gas
mixture.
27. When the reaction in one of the gas generators slows down so much that it is no longer useful, fill
the second gas generating tube with liquid (either HCl or H202, as appropriate) and use it instead.
Rocket Launches!
28. Collect the optimum (loudest) gas mixture one more time, and bring it to the instructor. Your
instructor will place the bulb on a rocket launch pad and ignite it with a piezo sparker. How far does
the micro mole rocket travel?
29. Collect the optimum mixture again, but this time leave about 1 ml, of water in the bulb. With your
instructor's consent, launch the micro mole rocket.
76
Name:____________________________
Class/Lab Period:___________________
Micro Mole Rockets
Data Table
"Pop-test" Properties of H2
Gas alone
"Pop-test" Properties of 02 Gas
alone
Pop-test Properties of 0 2:H2 Gas Mixtures
Oxygen:Hydrogen Mole
Ratio
1:5
Pop Test Results
2:4
3:3
4:2
5:1
Post-Lab Questions
1. In the space provided construct a bar graph to illustrate the pop test results with the various oxygen/hydrogen
gas mixtures.
77
2. Explain the relative loudness of pure oxygen and pure hydrogen in the pop-test.
3. Write a balanced chemical equation for the combustion reaction of hydrogen and
oxygen to give water.
4. When the reactants in a mixture are present in the exact mole ratio given by the
balanced chemical equation, all of the reactants should be used up when the reaction is
over. There will be no "leftover" reactants. However, if one of the reactants is present in
an amount greater than its mole ratio, then that reactant cannot react completely, and
some of it will be left over at the end of the reaction.
5. Use the mole ratio of hydrogen to oxygen from Question #4 to determine what happens
when various hydrogen/oxygen gas mixtures are allowed to burn. Complete the
following table to indicate which reactant (H2 or O2) is present in excess, and how
much of it will be left over after the combustion reaction is complete. Note: The
second one has been completed as an example.
Parts H2
6
5
4
3
2
1
0
Parts 02
0
1
2
3
4
5
6
Which reactant is present in excess?
How much of that reactant is left over?
H2
3
6. a) Which oxygen/hydrogen gas mixture produced the most explosive mixture? Provide
evidence to support your claim.
b) Explain why this mixture was most explosive.
7. Why do the hydrogen and oxygen gas mixtures in the collection bulb not react as soon
as they are collected? Note. Consider the role of the match and the properties of gas
molecules at room temperature.
78
http://www.csd509j.net/cvhs/kirscha/MicroMoleRockets%200708.doc
As adapted from Flinn ChemTopic- Labs - Molar Relationships & Stoichiometry
Micro Mole Rockets- Teacher Guide
Hydrogen and Oxygen Mole Ratio
Teacher Info taken from: http://teacherknowledge.wikispaces.com/Nelson+Inquiry+Lesson
Additional Info found in Flinn ChemTopic Labs “Molar Relationships & Stoichiometry” book, Vol. 7
Lesson Overview:
Students will learn how to use their knowledge of stoichiometry and mole ratios in order to
determine the best mixture of substances to propel a pipet rocket. The optimum ratio will be
used to calculate the mole ratio of the reaction of hydrogen and oxygen in a balanced chemical
equation. The concept of limiting reactants will be used to explain the results obtained with
various hydrogen/oxygen gas mixtures. Groups that have been already established earlier will
now compete in a friendly competition to see whose mole rocket can reach the furthest. This lab
is ultimately a culmination of the students’ knowledge of chemical equilibrium that has been
used in earlier lessons. This lab is a fun and informative lesson that helps to apply the subject
matter in a real-life situation that the students may not have realized existed. Students may not
have thought that mole ratios and chemical equilibrium play a major role in fuel, and allusions
can be made to automotives as well.
Learning Outcomes:
Students will understand mole ratios in a concrete fashion. Students will develop critical
thinking skills for lab procedures. Students will understand limiting reactants and will use that
knowledge to propel the micro mole rockets.
Michigan Standards:
C1.1 Scientific Inquiry, C4.6x Moles, C5.2X Balancing Equations, C5.2E Limiting Reactants
Students’ Prior Knowledge of Experiences:
Students should have already covered the topics that this lab reinforces; mainly how to determine
the limiting reactant, mole ratios, and the chemical equilibrium principles that these are built
from. The lab should come after the majority of the chemical equilibrium unit has been
discussed. Also, since this lab is most likely one to fall in the second semester, students should
have an understanding of common lab safety protocol, and general lab etiquette. Some common
issues that this lab should concern deal with student’s misconceptions of how mass relates to
mole ratio in a mixture, and also how to design a lab procedure that is based on sound reasoning.
Questions: (for use in class as well as a lab write-up, if desired)
1. How far, in meters, did your rocket fly? Do you believe it would be the same distance in outer
space? Why or why not? Explain your reasoning.
2. Why did you start by filling the rocket with water?
79
3. Which rocket would fly further?
a) a rocket filled with pure hydrogen
b) a rocket filled with a mixture of hydrogen and air
c) a rocket filled with a mixture of hydrogen and oxygen
4. Why must some water be left in the stem of the rocket in order for the launch to be
successful?
5. What is the reaction occurring inside the rocket?
6. Why is an electric charge needed? Explain this in relation to what you know about activation
energy.
7. Which hydrogen-oxygen rocket is expected to fly farther, a rocket that is mostly filled with
oxygen and some hydrogen or one mostly filled with hydrogen and some oxygen?
8. Would rockets filled with hydrogen and air fly at all?
9. What ratio of hydrogen to oxygen is optimal? Use a balanced equation to answer this
question.
10. Real rockets such as the NASA’s Saturn launch vehicle use liquefied gaseous hydrocarbons
and liquid oxygen for the rocket’s first stage and liquid hydrogen and liquid oxygen for the
second stage. What sort of design feature would keep liquid hydrogen and liquid oxygen from
reacting until they are supposed to?
Cautions:
The ignition and launching of the rockets should be closely monitored for student’s safety.
Goggles must be worn throughout the experiment, including during the launching. For extra
precaution, a glove can be worn for ignition. All other general lab safety protocols should be
followed as well. Set aside an area for launching the rockets. Set up a target and do not let
students in that area while rockets are being launched.
Materials needed to construct the Piezoelectric Sparking Mechanism:
1. piezoelectric sparking mechanism (modified Charcoal Lighter such as the Aim ‘n Flame Torch
lighter)
2. electric tape
3. 6-D common nail
4. film canister
5. hot glue gun and glue
6. 10 cm of 24 AWG speaker hook-up wire (double solid, not stranded: Radio Shack #2781509)
7. wire strippers
Constructing the Piezoelectric Sparking Mechanism:
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1. Obtain a piezoelectric lighter from a hardware store. Modify by attaching a wire to the end to
separate the sparking cap.
2. Using scissors split the end of the 10-cm length of speaker wire down the center for a distance
of 3-4 cm. Then strip the last 1-2 cm of insulation from both sides of the “fork”.
3. Drill a 1/8” hole (or alternately, melt a hole with a hot 6D common nail) through the center of
the film canister lid. Slide the wire through, and secure the lid in place about 4 cm from the
other end (end without a fork) of the wire with some hot melt glue, or other water-proof
adhesive. The film canister serves both as a launch support pad and as a splash guard.
4. Slide one of the two stripped wire ends down inside the hold in the end of the charcoal lighter.
It should be inserted all the way down so that the insulation abuts the nozzle. (It helps to rotate
the wire as you insert it.) Lay the other stripped end alongside the metal shaft of the lighter and
secure it in place with 3-4 wrappings of electric tape. A water shield is a necessity for this
experiment. Check the hot glue seal on the film canister lid. Test the device by pulling the
trigger and looking for a small spark to flash across the far end of the wire.
5. Note: If no spark appears, there could be a short occurring through the insulation if the wire
that extends from the nozzle comes too close to the rim of the shaft. If so, centering the wire and
surrounding it with insulation such as hot melt glue or a few wrappings of electric tape, should
solve the problem.
Another Set of Questions and Answers:
Q1. Write balanced equations for the reactions taking place inside the two generator vials.
A1. 2 HCl(aq) + Zn(s) --> ZnCl2(aq) + H2(g)
Note: we do not see the zinc chloride that is produced along with the hydrogen gas, for the zinc
chloride is soluble and therefore remains as dissociated ions in solution.
2 H2O2(aq) --> 2 H2O(l) + O2(g)
Note: likewise in this reaction, the by-product (water) is not evident, for it simply mixes in with
the water already present as a solvent for the hydrogen peroxide.
Q2. On the graph attached draw a bar graph for the relative loudness produced by pop-testing the
various mixtures. Use the right-hand graph for plotting class averages.
81
A2.
Q3. Explain your observations for the pop-test of pure H2 and of pure O2.
A3. Hydrogen is combustible; this means it reacts readily with oxygen. Pure hydrogen therefore
is not by itself an explosive substance. If a slight pop is heard, it was due to the presence of
oxygen (usually from air left in the generator). Oxygen shows a negative pop-test. Oxygen
supports combustion but is not itself combustible.
Q4. What proportion (H2 : O2) produced the most explosive mixture? Why was that mixture
most explosive? (Hint: Write a balanced equation for the reaction of H2 and O2.)
A4. The most explosive mixture of hydrogen to oxygen is 4:2, or simply 2:1. This is the mole
ratio from the balanced equation (see equation). It is most explosive because it allows for the
maximum yield of product (water) and the maximum output of heat. Because the two reactants
are present in this optimum ratio, they are both completely consumed; in other words, nothing is
wasted.
2 H2(g) + O2(g) --> 2 H2O(g) + Q
Q5. Why do the H2 and O2 in the collection bulb not react as soon as they mix? What role does
the spark play?
A5. Even if H2 and O2 are both present in a combustible ratio, and the H2-O2 collisions are
occurring at a considerable rate, the collisions are generally not occurring with enough energy to
form the activated complex, and the reaction can not proceed at a detectable rate. By supplying
extra energy, in the form of heat or electricity, the particles move faster and collide harder on
average and therefore have a greater chance of forming the activated complex and enabling the
reaction to occur at reasonable rate. This minimum energy requirement is known as the
activation energy for the reaction.
Q6. What methods did you attempt for making your rocket fly farther? Which ones worked?
A6. Several parameters influence the length of the rocket's flight. One of course is the ratio of the
gases, for reasoning already mentioned. A second is the angle of the launch: 45º is best under
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ideal, frictionless conditions, but with the considerable air resistance, some angle less than 45º
invariably proves more effective. The mass of the rocket plays a major role, a weighted rocket is
less subject to air resistance, but it also has more inertia to overcome. Streamlining the rocket
(with tail-fins and a nose-cone, for example) can also increase its flight. Leaving some water in
the bulb can greatly increase the flight as well, for it gives the expanding gases something to
push against, a propellant as it were. One can find a good illustration of this principle in toy
water rockets, which rely on pumping air into a plastic bulb partly filled with water.
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Moles & Stoichiometry
Activity #10 – Explosive Stoichiometry
Questions to be investigated
What are the main principles of stoichiometry? Why is a balanced equation
important when performing calculations and labs? What specific amount of a
particular gas is needed to fill the tennis ball canister to produce the best
explosion?
Objectives
The student will understand stoichiometric conversions and balanced equations.
The student will be able to calculate the exact amount of a particular gas that is
needed to fill a tennis ball canister to produce the best explosion in this lab.
Teacher Notes
See attached page called “CLIFF NOTES FOR THE TEACHER”. This lab should
only be done under careful supervision close to the end of the unit. It could be
used as an assessment at the end of a unit. You should practice this lab on your
own before using it in a classroom. It produces a VERY loud noise, so check
with administration before performing this lab in the building. It’s a cool lab, but
if you have any doubt about your student’s ability to listen and follow directions,
think twice.
Materials
Goggles, Aprons, Ignitor, 60 mL syringe, 850 mL Tennis ball canister, Methane
gas, Butane gas, Propane gas
Safety Concerns
Methane, butane, and propane gases are highly flammable; the explosion
resulting from the ignition of the three gases will be VERY loud. Check with
administration before performing this lab. Hearing protection is advised, such as
ear plugs or ear muffs.
Real-World Connections
Masses and volumes of products are extremely important in key reactions such
as the deployment of air bags in vehicles, carbon dioxide production to help
bread rise, and the production of ammonia gas for industries.
Sources
http://moodle.oakland.k12.mi.us/clarenceville/file.php/47/Explosive_Stoichiometr
y_Lab.doc
(This is the website for Craig Reisen from Clarenceville High School. There may
be a key needed to log in to the site, but if you email Craig, he will give it to you.
This activity is listed under Chemistry B on the Courses page.)
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Procedure/Description of Lesson
See next page for student lab sheets and teacher notes.
Assessment Ideas
This lab could be used as an assessment at the end of a unit. Points or a grade
could be given based on the closeness to the correct answer.
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Name ______________________________
Date _________
Lab Sheet
Hour _________
Explosive Stoichiometry
Purpose:
To investigate some principles of stoichiometry, writing and balancing
equations, and the ideal gas law to create a controlled explosion inside of a
tennis ball canister.
Discussion: You will need to refer to your notes on the Ideal Gas Law in this lab.
Almost all organic molecules and some inorganic compounds will burn in air. Many will ignite
quite readily and are considered flammable. When a compound or molecule burns in air, it is
actually reacting with the oxygen in the air. As a result, chemists refer to combustion reactions
as oxidation reactions. CH4 + 2 O2 Æ CO2 + 2 H2O is an example. The general form of
a combustion reaction is:
hydrocarbon + oxygen Æ carbon dioxide + water.
In this lab, you will have to calculate the exact amount of a particular gas needed to fill a tennis
ball canister in order to produce the best explosion. If your calculation is low or high, no
explosion will occur.
Materials
Goggles, Aprons
Ignitor
60 ml Syringe
850 ml Tennis Ball Canister
Methane Gas
Butane Gas
CAUTION: This lab may be very hazardous to your health!
Procedure
1. Go to the Calculations and Data Section and complete all the necessary calculations BEFORE going
on to perform the actual lab.
2. Have the teacher check your calculations when completed. YOU ONLY GET ONE TRY AT THIS
for each gas! We may only have time for each group to do one gas.
3. Obtain a pair of goggles, an apron and sheet of cardboard for your lab station.
4. Go to the teacher to get the sample of gas (butane or methane).
a. Fill the 60 ml syringe to the appropriate volume based on your calculations.
b. Transfer the gas from the syringe to the sealed tennis ball canister through the small
hole just above the bottom of the canister.
c. Give the system 10-15 seconds of wait time to allow the gas to diffuse.
d. Insert the ignitor.
e. Cover your ears and yell “FIRE IN THE HOLE.” Press the ignitor button.
f. If you are successful, have the teacher initial the box to the right.
g. If you are unsuccessful, open the top of the tennis ball canister and vent it while
you recalculate and try the other gas.
5. Clean up.
Teacher
initials
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6. Calculations and Data
measurements.
Show all work and units for ALL calculations &
1. Write down the formula of the gas that you will be using: _____________________.
2. Write a balanced chemical equation below for the combustion of that gas:
3. Record the current temperature in the room: __________ … ___________.
4. Record the current atmospheric pressure in the room: ______________.
5. The canister is 850 ml and air is 21% oxygen. Calculate the number of milliliters of oxygen
that are inside the canister when it is sealed.
6. Calculate the number of moles of oxygen that are in your canister.
7. Using your balanced chemical equation, calculate the number of moles of your gas that you
will need to react completely with all of the oxygen that is in the canister.
8. Using your previous answer and the current temperature and atmospheric pressure, calculate
the number of milliliters of your gas that you need to add to the canister. (pay attention to
units).
Conclusions and Questions
1. Is the combustion reaction in this lab endothermic or exothermic? How do you know?
2. Based on your answer to question #1, what happens to the volume of the gas as the reaction
progresses? Does this help or hinder the desired outcome of pooping the lid off the canister?
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CLIFF NOTES FOR THE TEACHER
Methane
•
•
•
•
•
•
•
CH4 + 2 O2 Æ CO2 + 2 H2O
850 ml x 21% = 178 ml = 0.178 L of O2 (g) in the tennis ball canister
0.178 L x 22.4 L/ mole = 0.008 moles of O2 (g)
Using mole ratios Æ 0.004 moles of methane (CH4)
0.004 moles x 22.4 L/mole = 0.089 L = 89 ml
89 ml of methane gas needed to fill the tennis ball canister
Students will need to inject the syringe twice to get the 89 ml into the tennis ball
canister
Butane
•
•
•
•
•
•
•
2 C4H10 + 13 O2 Æ 8 CO2 + 10 H2O
850 ml x 21% = 178 ml = 0.178 L of O2 (g) in the tennis ball canister
0.178 L x 22.4 L/ mole = 0.008 moles of O2 (g)
Using mole ratios Æ 0.0012 moles of butane (C4H10)
0.0012 moles x 22.4 L/mole = 0.0266 L = 27 ml
27 ml
Containers of Butane can be purchased at a dollar store Æ find containers that have a
small “neck” to dispense the gas into the syringe.
Propane
•
•
36 ml
Propane burners (for soldering, etc.) work, but you have to jury-rig a plastic dropper
pipette to seal with the syringe
•
Weather.com will give the prevailing atmospheric pressure
•
Students need to convert to K temperature when using the ideal gas law
•
A Bar-B-Que gas grill igniter is ideal for the igniter system needed in this lab
Lid
Tennis ball
Small opening to allow
igniter to fit inside
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