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Unit 5: Moles and Chemical Reactions Notes
Unit Objectives:

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Understand the mole concept
Calculate gram formula mass
Differentiate between formula mass and gram formula mass
Convert between grams and moles
Balance a chemical reaction by adjusting only the coefficients
State the Law of Conservation of Mass and Energy and relate it to balanced chemical equations
Create and use models of particles to demonstrate balanced equations
Identify various types of reactions: synthesis, decomposition, single replacement, & double replacement
Solve mole-mole Stoichiometry problems given a balanced reaction
Calculate the empirical formula from percent mass
Differentiate between empirical and molecular formulas
Determine the molecular formula from the empirical formula and molecular mass
Define the following vocabulary:
Empirical formula
Mole
Percent Composition
Formula mass (FM)
Subscript
Molar Mass
Reactant
Gram formula mass (GFM)
Product
Coefficient
Law of conservation of mass and energy
Balanced equation
Synthesis reaction
Decomposition reaction
Single-replacement reaction
Double-replacement reaction
Molecular formula
1
The bold, underlined words are important vocabulary words that you should be able to define and use properly in
explanations. This is a study guide for what you will be tested on throughout the year. The objectives are divided into
categories of “Knowledge” (what you have to know) and “Application” (what you have to be able to do).
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I. CHEMICAL FORMULAS, REACTIONS & STOICHIOMETRY
Knowledge
Application
Chemical formulas are used to represent compounds.
The main types of chemical formulas include: empirical,
molecular, and structural.
o Determine the empirical
An empirical formula is the simplest whole-number ratio of atoms
formula from a molecular
in a compound.
formula
Molecular formulas are chemical formulas that show the actual ratio
o Draw structural formulas for
of atoms in a molecule of that compound.
covalent (molecular)
Structural formulas can also be used to represent covalent
compounds
compounds. These use lines to show covalent bonds between atoms
and also show the geometrical arrangement of atoms in the
compound.
One mole of any substance is equal to 6.02 x 1023 pieces of that
o Calculate the molar mass (gramsubstance.
formula mass) of a substance
The formula mass of a compound is equal to the sum of the atomic o Determine the molecular
masses of its atoms (units are atomic mass units)
formula, given the empirical
The molar mass (gram-formula mass) of a substance is equal to
formula and the molar mass
the formula mass in grams – hence “gram-formula mass”.
o Determine the number of moles
The mass of one mole of any substance is equal to its molar mass
of a substance, given its mass
(gram-formula mass).
and vice versa
o Calculate the percent
composition of any element in a
The percent composition by mass of each element in a compound
given compound
can be calculated mathematically.
o Calculate the percent
composition of water in a given
hydrate
o Balance equations, given the
Balanced chemical equations show conservation of matter, energy,
formulas for reactants and
and charge.
products
The coefficients in a balanced equation can be used to determine
o Calculate simple mole-mole
mole ratios in the reaction.
ratios, given balanced equations
o Identify the different types of
Types of chemical reactions include synthesis, decomposition,
chemical reactions, given their
single replacement, and double replacement.
chemical equations
2
Mini Poster Project- Classifying Reaction Types
Background
There are many different types of chemical reactions in the world. We have studied five
types of them during this unit. You are now going to create a mini poster describing each
chemical reaction and differentiating between the types of chemical reactions.
This will be an independent project. During your allotted time you will further research
each of the types of reactions as well as find real-world examples and pictures of each
reaction type for visual aide on your poster.
Objective
In this project you will create a mini poster that provides information about the five types
of chemical reactions (synthesis, decomposition, single replacement, double replacement
and combustion). This posted should be created on a manila folder as a bi-fold or as a trifold.
Materials
You will be provided with colored pencils, markers, paper, and folders. If there is
anything additional that you would like to use on your poster you must supply it.
Procedure
The following information should be included for each reaction type:
• Name of reaction type
• Generic equation
• Real Life example of reaction
• Definition of reaction type
Each
•
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bullet point below should also be considered:
Title of Poster
Color
Pictures
Readability
Sizing & Usage of space
Creativity
Originality
3
Lesson 1: The Mole & Gram Formula Mass(GFM)
Objective: To define and calculate molar mass. To apply the formula relating mass in grams to moles
Chemistry is a basic science whose central concerns are 
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the structure and behavior of atoms (elements)
the composition and properties of compounds
the reactions between substances with their accompanying energy exchange
The laws that unite these phenomena into a comprehensive system.
We have already studied the structure and behavior of atoms (elements) as well as the composition and properties of
compounds. We will now be moving on to learn about the reactions that occur between substances. In order to do this
we must be able to quantify the reactants and products to see the changes that are occurring.
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What is a Mole? = 6.02 x 1023 things
 Symbol:
 Like a dozen, only waaaaaaaaaaay bigger
 Why we use it: atoms are soooooo tiny that we have LOTS of them in a given sample
1 dozen eggs = 12 eggs
1 mol eggs = 6.02 x 1023 eggs
Both subscripts and coefficients represent MOLES
Ex:
in 1 mole of NaCl, there are
1 mol of Na+ ions
1 mol of Cl– ions
in 1 mole of H2O, there are
2 mol of H atoms
1 mol of O atoms
in 1 mole of Na2SO4, there are 2 mol of Na+ ions
1 mol of SO4 2- ions
How many moles of atoms in total?
2 mol of NH4 + ions
1 mol of CO3 2- ions
in 1 mole of (NH4)2CO3, there are
How many moles of atoms in total?
*Honors Chemistry Only:
*Hydrates:
Ex: Na2CO3  7 H2O
7
in 1 mole of this substance, there are
Molar Mass:
a.
Example: the molar mass of carbon is 12.011 g/mol
b. molar mass unit is :
c. molar mass is also called :
d.
Example: H2O =
5
14
Classwork 5-1:
1. Fill in the table below. Put an “M” if the substance is molecular/covalent and an “I” if ionic.
Formula
Example: HClO3
a.
M
NH4C2H3O2
Moles of each atom
1 mol H atoms
1 mol Cl atoms
3 mol O atoms
Gram
formula
mass
Formula
Moles of each atom
CaCl2
84
g/mol
d.
Mg3(PO4)2
b.
e.
Mg(OH)2
CH3CH2CH3
c.
f.
6
Total
moles of
atoms
Chemical Formula
1. MgBr 2
Gram Formula M ass
2. KCl
3. FeCl 2
4. CrF 2
5. Al 2 S 3
6. PbO
7. TiI 4
8. Mg 3 P 2
9. SnCl 2
10. HgCl 2
11. H 2 SO 4
7
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Lesson 2: Mole-Mass Conversions
Summary: Gram-formula mass (also called “molar mass” or GFM)
Definition:
Calculate:
Unit:
Ex:
Calculate the gram-formula mass
of NaCl
Calculate the gram-formula mass
of Ba(NO3)2
# Atoms x Mass
# Atoms x Mass
Na
Cl
Ba
N
O
We will need to convert from grams to moles and vice versa for this class. The diagram below summarizes these
processes:
Converting Grams to Moles:

From Table T, you would use the MOLE FORMULA:
Example: How many moles are in 4.75 g of sodium hydroxide? (NaOH)

Step 1: Calculate the GFM for the compound.
Na = 1 x =
O =1x=
H =1x=

Step 2:
9
=
=
= +
=
Converting Moles to Grams:
The same formula can be used to convert moles back to grams
Example: You have a 2.50 mole sample of sulfuric acid (H2SO4). What is the mass of your sample in grams?

Step 1: Calculate the GFM for the compound.
H = 2 x 1.0= 2.0
S = 1 x 32.0= 32.0
O= 4 x 16.0= 64.0
= 98.0 g/mol

Step 2: Plug the given value of moles and the GFM into the “mole calculation” formula and solve for the mass of
the sample.
grams =
= 245 g of H2SO4
Classwork 5-2:
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Classwork 5-2(con't):
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12
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Activity - Moles of Snacks
Directions: Take the necessary measurements, and record them with units and without rounding.
Show all your calculations, rounding your answers to three significant digits and labeling units
clearly. Remember to round molar mass to the tenth (0.1) place. Show ALL work!
Pre-lab
1. Find and record the molar mass of sodium________________
2. Find and record the molar mass of sodium chloride_____________
Materials
Small bag of snack food
Balance
Procedure
1. Using the information on the snack package answer the following questions.
SHow ALL work for any conversions that must be done Final answers must be in three
significant figures. Molar masses must be rounded to the tenth place with units.

What is the mass of one serving of the snack in the snack bag?
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How many mg of sodium are in one serving of the snack?
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How many g of sodium are in one serving of the snack?
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How many moles of sodium are in one serving of the snack?
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How many individual atoms of sodium are in one serving of the snack?
2. Using the electronic balance, mass the bag of snack food and record the result__________
3. Open the snack bag and either empty the contents in the designated container in the lab or
eat the contents, sharing with the entire lab group.
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Name:_____________________________________________________Date:_________Period:______
4. Using the electronic balance, mass the empty snack bag and record the
result_________________
5. Calculate the mass of the snack food that was in the bag and record the result. Show your
work below. Circle your final answer.
6. Is the mass you calculated in #5 the same as the mass of one serving that was recorded on
the label of you snack bag. Yes/No (circle one). Why do you believe this happened? ________
______________________________________________________________________________
______________________________________________________________________________
______________________________________________________________________________
______________________________________________________________________________
______________________________________________________________________________
______________________________________________________________________________
______________________________________________________________________________
Extension: Healthy American adults should restrict their sodium intake to no more than 2,400
milligrams per day. This is about 1¼ teaspoons of table salt (sodium chloride [NaCl]).
Answer the following questions.
SHow ALL work for any conversions that must be done Final answers must be in three
significant figures. Molar masses must be rounded to the tenth place with units.
1. What is the maximum number of moles of sodium recommended in your diet? How many
sodium atoms would this be?
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Moles Lab Activity –Aluminum
Remember: Take the necessary measurements and record them with units. Do not round these
numbers. Show all calculations. Molar mass is rounded to the tenth (10th) place. All other calculations
are rounded to three (3) significant figures. Include units in your final answers.
Materials
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Empty aluminum can
Balance
Aluminum foil
Procedure
1. Determine the molar mass of aluminum and record it here:__________________
2. Place the aluminum can on the balance and record the mass._____________
3. Answer the following questions:

Does the aluminum can contain more than, less than, or exactly one mole of aluminum? Explain
your answer using complete sentences or calculations.

Calculate how many moles of aluminum are in one aluminum can. Show ALL work!!

Calculate how many atoms of aluminum are in one aluminum can.
Extension
1. How many cans would you need to have one mole (1.00 mole) of aluminum? Show your work using
the factor label method.
2. Aluminum foil is essentially pure aluminum that has been pounded into a thin sheet. How many
grams of aluminum foil would be necessary to have 1.00 mole of aluminum foil? Explain your
answer using complete sentences or calculations. Have your teacher check your answer before
going on to the next step.
3. Go to the balance and measure out exactly one mole (1.00 mole) of aluminum foil. As a lab group
make a “sculpture” out of the aluminum foil, using the entire mole of foil.
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Lesson 3: Types of Formulas
Objective: To compare and contrast Empirical and molecular formulas. To calculate the molecular formula from the
empirical formula and molecular mass.
In order to make it easier to describe elements and molecules, chemical formulas are used.
-
The empirical formula is the
***Ionic compounds are already empirical formulas.
The molecular formula is the
It is always whole number multiples of empirical formula. The Subscripts can be reduced same as when you
reduce a fraction. Remember…subscripts are just ratios of how the atoms are coming together in a
compound!
Determining Empirical Formula
 Divide subscripts by the greatest common factor
Example: molecular formula = C4H10

Divide by 2 (greatest common factor)
Answer:
C2H5
Determining MOLECULAR Formula from the Empirical formula and molecular mass

Step 1: Calculate Gram Formula Mass of the EMPIRICAL FORMULA

Step 2: Divide the MASS of the molecular formula by the MASS of the empirical formula.

Step 3: MULTIPLY the subscripts of the empirical formula by the answer in step 2.
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Example: The empirical formula of a compound is C2H3, and the molecular mass is 54.0 grams/mole. What is the
molecular formula?
Step 1. Determine the GFM of the empirical formula.
C2H3 = (2 C X 12.0 g/mol) + (3 H X 1.0 g/mol) = 27.0 g/mole
Step 2. Divide the molecular mass by the empirical mass. This will give you a whole-number multiple that tells you how
many times larger the molecular formula is than the empirical formula.
(54.0 g/mol) / (27.0 g/mol) = 2
Step 3. Multiply the whole number by the empirical formula. This will give the molecular formula.
2 X C2H3 = C4H6
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Classwork 5-3:
1. Identify each of the following as an empirical or molecular formula. If a formula is molecular, write its empirical
formula.
Formula
Empirical or
Simplify if
Molecular?
Molecular
Formula
NaCl
N2O4
C2H6
Ra(CN)2
C6H12O6
Ba(NO3)2
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Empirical or
Simplify if
Molecular?
Molecular
6. What is the molecular formula of a compound that has a mass of 289g and an empirical formula of NH3?
7. What is the molecular formula of a compound with a mass of 760g and an empirical formula of Cr 2O3?
8. What is the molecular formula of a compound that has an empirical formula of NO 2 and molecular mass of 92.0 g?
9. A compound has an empirical formula of HCO2 and a molecular mass of 90 grams per mole. What is the molecular
formula of this compound?
10.VitaminChasanempiricalformulaofC3H4O3andamolecularmassof264.0g/mole.
Determinethemolecularformula.
11.Ibuprofen,acommonheadacheremedy,hasanempiricalformulaofC7H9Oandamolar
massof545.0g/mole.Determinethemolecularformula.
12. OxalicacidhastheempiricalformulaHCO2andamolarmassof90.0g/mole.
Determinethemolecularformula.
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Methane Bubbles Demo
Prelab
1.
2.
3.
Questions
What does the word combustion mean?
What type of gas is methane and what is its chemical formula?
What do you thing will happen if we added fire/heat to methane gas?
Analysis
1. Why are the methane bubbles able to rise?
2. Write the balanced chemical equation for the combustion of methane.
methane reacts with oxygen to produce carbon dioxide, water, and heat)
3. What are the products for the complete combustion of all hydrocarbons?
4. List three examples of a hydrocarbon other than methane.
5. What do you think will happen if you placed some of the methane bubble
solution on your hands and touched a burning candle to it?
6. Why is your hand unharmed by the flame?
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Lesson 4 Types of Chemical Reactions
Objective: Identify various types of reactions: synthesis, decomposition, single replacement, & double replacement
1. Synthesis
During a
General pattern:
Example:
2Mg(s) + O2(g) → 2MgO(s)
2. Decomposition
During a
General pattern:
Example:
2HgO(s) → 2Hg(ℓ) + O2(g)
3. Single Replacement
During a
General Pattern:
Use Table J
Example:
Zn(s) + CuSO4(aq)
Cu(s) + ZnSO4(aq)
Cu(s) + ZnSO4(aq)
4. Double Replacement
During a
General Pattern:
Example: AgNO3(aq) + NaCl(aq) → NaNO3(aq) + AgCl(s)
A PRECIPITATE
Check Table F to see if there is a PRECIPITATE during a double replacement reaction.
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Classwork 5-4:
Practice: Determine if the following reactions are synthesis (S), decomposition (D), single replacement (SR), or double
replacement (DR) reactions.
1. H2
+
Cl2

2HCl
________
2. NaOH
+
HCl

NaCl
3. 2NH3

N2
+
3H2________
4. Mg
+
H2SO4
5. F2
+
2HBr
6. Zn(NO3)2
+

+
H2O________
MgSO4 +
H2________

+

CaCO3
Br2
Ca(NO3)2
3O2

2Al2O3________
8. CuSO45H2O

CuO
+
9. SiF6
6Xe

SiXe6
7. 4Al
+
+
2HF________
H2SO4
+
10.
2 NaClO3 2 NaCl + 3 O2 ________
11.
2 AgNO3 + Ni Ni(NO3)2 + 2 Ag________
+
+
4H2O________
3F2________
12. H2CO3 H2O + CO2 ________
13.
BaCO3 BaO + CO2________
14.
4 Cr + 3 O22 Cr2O3________
15.
Ca + 2 HCl CaCl2 + H2 ________
16.
Ca(C2H3O2)2 + Na2CO3 CaCO3 + 2 NaC2H3O2 ________
17.
Cu(OH)2 + 2 HC2H3O2 Cu(C2H3O2)2 + 2 H2O________
18.
8 Cu + S8 8 CuS ________
19.
P4 + 5 O2 2 P2O5________
20.
2 K + 2 H2O2 KOH + H2 ________
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ZnCO3________
Single Replacement Lab
Objective: Learn how to use the activity series to predict the products of a single replacement
reaction.
Materials:
6M HCl
Well Plates
Scoops
Metals (magnesium, iron, zinc, copper, aluminum)
Procedure:
1. Label each well for a particular piece of metal on a separate
Mg Cu Fe Al Zn
piece of paper.
2. Put a very small piece of the metal into each well as
assigned in step #1.
3. Place 2 drops of the 6M hydrochloric acid into the first well. Observe the reaction.
4. Repeat step 4 for each of the other metal samples.
5. Rinse off larger pieces of metal before disposing of them in the trash can. Rinse all
wells of the tray. Let dry.
Observations: Note any observations for each metal
Mg
Cu
Fe
Al
Zn
Analysis:
Write a balanced chemical equation for each reaction that occurred. Include the
proper states of matter (i.e. (s) = solid, (l) = liquid, (g) = gas, (aq) = aqueous
solution).
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Lab - Single Replacement Reactions & Table J
Introduction
Not all chemicals react when combined. The activity series is a listing of the level of reactivity that elements possess. It will help
you determine whether or not a specific single replacement reaction will occur.
Purpose
The purpose of this demonstration is to exhibit how the Activity Series can be useful in determining whether or not a single
replacement reaction will occur.
Predictions
The following “reactants” (chemicals) will be mixed together. I hesitate to call them all reactants because they will not all react with
each other. Determine which will react and which will not using the Activity Series.
Will a reaction occur?
Did a reaction occur?
#
“Reactants” (Chemicals)
1
CuSO4 + Fe →
2
Al + CuSO4 →
3
NaCl + Al →
4
Mg + CaCl2 →
5
NaCl + Cu →
6
HCl + Al →
Yes
No
Follow-up
1.
After the demonstration, indicate in the table above whether or not a reaction took place (in the furthest-right column).
2.
For the reactions that took place:
a. Finish the chemical equations in the table above by writing what products were formed.
b. Balance each of the reactions in the space below. Be neat.
3.
For the reactions that did not take place:
Finish the chemical equations in the table above by writing ‘NR’ on the product side. This indicates that no reaction will
take place.
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Lesson 5: Balancing Equations
Objective: Balance a chemical reaction by adjusting only the coefficients
Review:

REACTANTS =

PRODUCTS =

COEFFICIENT =

SUBSCRIPT =
*COEFFICIENTS and SUBSCRIPTS tell us how many moles we have for each element
Chemical Symbols (states of matter)
(s)
(g)
(l)
(aq)
l

Example: : 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
CATALYST=
Example: 2 H2O2(aq) ---- -> 2 H2O(l) + O2(g)
The catalyst in this example is KI
Law of Conservation of Mass
A chemical equation is a set of symbols that state the products and reactants in a chemical reaction.
This goes for mole amounts as in balancing equations and mass amounts.
Example: Given the reaction: N2 + 3H2 ---------- 2NH3
What is the total number of grams of H2 that reacts when 14 grams of N2 are completely consumed to
produce 20 grams of NH3?
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An equation is balanced when:


**NOTE:
Steps for Balancing Equations
Step 1: Find the most complex compound in the equation. Balance the elements found in that compound on the
opposite side of the arrow by changing the coefficients for those atoms.
Step 2: Continue balancing until all atoms are balanced (save pure elements for last)
Step 3: Go back and check each atom to see if it is balanced on both sides of the equation.
Step 4: POLYATOMIC IONS may be balanced as a SINGLE UNIT rather than as separate elements as long as they stay
intact during the reaction.
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Lab -Investigating the Law of Conservation of Mass
Objective: To corroborate the law of conservation of mass through laboratory experimentation.
Background:
When wood burns or water evaporates, some or all of it appears to disappear but does it? A pile of
ash is a lot smaller than the wood it originally started as and a dry empty cup seems a lot lighter if all
of the water is allowed to evaporate. When you pop open a soda can, the fizz is really carbon
dioxide, tiny gas bubbles that rise out of the soda and mix back with the atmosphere. To the naked
eye, it seems like these materials just vanish. However, does matter ever really disappear?
According to the law of conservation of mass, mass is always conserved. In many situations, the end
products appear to have decreased but when all the products are captured and their mass
measured, it was discovered that the mass of the reactants is equal to the mass of the products. In
other words, matter can neither be created nor destroyed. This is called the law of conservation of
mass. All atoms present in the reactants are also present in the products.
The "Law of Conservation of Mass" states that when matter goes through a physical or
chemical change, the amount of matter stays the same before and after the changes occur.
In other words, matter cannot be created or destroyed.
According to this principle, matter cannot just appear or disappear, but matter can change its
nature. When atoms chemically combine with or split away from other atoms the chemical and
physical properties of that substance can be completely altered. Take table salt, for example. The
chemical formula for salt is NaCl, or sodium chloride. Sodium by itself is an element an unstable and
poisonous metal. Chlorine by itself is also an element, which at room temperature is a poisonous
gas. However, combine these two atoms together with a chemical bond and they are completely
transformed into a very stable substance that is edible crystalline solid at room temperature.
Materials:
Sodium bicarbonate, NaHCO3
scoopula
Funnel
Acetic acid, CH3COOH
250 mL Erlenmeyer flask
Balloon
Triple beam balance
Graduated cylinder
Masking tape
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Procedure:
1.
Observe the physical properties of acetic acid and sodium bicarbonate and record your
observations in data table 1.
2.
Put your funnel inside the balloon. Make sure the neck of the funnel is deep inside the
balloon, so you do not get any sodium bicarbonate stuck in the neck of the balloon.
3.
Put 4 scoops of sodium bicarbonate in the funnel.
4.
Measure out 50 mL of acetic acid. Pour the acetic acid into the Erlenmeyer flask.
5.
Stretch the balloon over the mouth of the flask, with the balloon hanging to the side of the
flask. (DO NOT let the reactants mix). Use masking tape to seal the balloon to the neck of
the flask.
6.
Record the combined mass of the ENTIRE SETUP in data table 2
7.
Lift the balloon so that the sodium bicarbonate now drops into the Erlenmeyer flask.
8.
Observe for 5 minutes.
9.
Observe the physical properties of the products and record your observations in data table
1.
10.
Record the mass of the ENTIRE SETUP again data table 2.
Data:
Table 1:
Substance
reactants
Acetic Acid,
CH3COOH(aq)
Sodium Bicarbonate,
NaHCO3(s)
products
29
NaCH2COOH(aq)
H2O(l) +CO2(g)
Physical Properties
Table 2:
Total Mass Before Reaction
(grams)
Total Mass After Reaction
(grams)
Results:
1.
What evidence that a chemical reaction took place did you observe when the sodium
bicarbonate and acetic acid were mixed?
2.
Calculate the change in mass that you measured between the initial reactants and final
products of the reaction.
3.
Calculate the percent error between your measured final mass and the final mass predicted by
the law of conservation of mass.
4.
Why was it necessary to do this experiment in a sealed container (closed system) in order to
accurately illustrate the principle of conservation of mass?
30
Classwork 5-5:
LawofConservationofMassWorksheet
Since there is a conservation of mass in all balanced chemical reactions, these problems are also quite easy.
Consider the following balanced chemical reaction CaCO3 à CaO + CO2. In this reaction the mass of
CaCO3 must equal the combined mass of CaO and CO2. If you know the mass of the reactants you
automatically know the mass of the products and vice-versa.
Using the reaction above, determine the mass of CaO produced if 200 grams of CaCO3 decomposed and
produced 88 grams of CO2.
CaCO3 à CaO + CO2
200g = Xg + 88g à 200g – 88g = 112g CaO produced
Complete the following problems
1) Hydrogenandoxygenreactchemicallytoformwater.Howmuchwaterwouldformif14.8grams
ofhydrogenreactedwith34.8gramsofoxygen?(H2+O2àH2O)
2) Whenammoniumnitrate(NH4NO3)explodes,theproductsarenitrogen,oxygen,andwater.When
40gramsofammoniumnitrateexplode,14gramsofnitrogenand8gramsofoxygenform.How
manygramsofwaterform?(NH4NO3àN2+O2+H2O)
3) 40gofcalciumreactswith71gofchlorinetoproduce_____gofcalciumchloride.
4) _____gofpotassiumreactswith16gofoxygentoproduce94gofpotassiumoxide.
5) 14goflithiumreactionwith_____gsulfurtoproduce46goflithiumsulfide.
6) 24gofmagnesiumreactswith38goffluorinetoproduce_____gmagnesiumfluoride.
7) 65.5gcopperreactswith_____goxygentoproduce81gcopper(I)oxide.
8) 88gofstrontiumreactswith160gbrominetoproduce_____gstrontiumbromide.
9) 46gofsodiumreactswith_____gchlorinetoproduce117gsodiumchloride.
10) ____gironreactswith71gchlorinetoproduce129gofiron(II)chloride.
11) 137gofbariumreactswith______giodinetoproduce391gbariumiodide.
12)_____ghydrogenreactswith32gofoxygentoproduce34gofhydrogenperoxide.
13)Whydowebalancechemicalreactions?
31
Practice : Balance the following equations in the space provided - Use a pencil
1)
___ NaNO3 + ___ PbO  ___ Pb(NO3)2 + ___ Na2O
2)
___ AgI + ___ Fe2(CO3)3  ___ FeI3 + ___ Ag2CO3
3)
___ C2H4O2 + ___ O2  ___ CO2 + ___ H2O
4)
___ ZnSO4 + ___ Li2CO3  ___ ZnCO3 + ___ Li2SO4
5)
___ V2O5 + ___ CaS  ___ CaO + ___ V2S5
6)
___ Mn(NO2)2 + ___ BeCl2  ___ Be(NO2)2 + ___ MnCl2
7)
___ AgBr + ___ GaPO4  ___ Ag3PO4 + ___ GaBr3
8)
___ H2SO4 + ___ B(OH)3  __ B2(SO4)3 + ___ H2O
9)
___ Fe2O3 + ___ H2  ___ Fe + ___ H2O
10)
___ Li + ___ N2  ___ Li3N
11)
___ Zn + ___ HCl  ___ ZnCl2 + ___ H2
12)
___ NaCl + ___ AgNO3  ___ NaNO3 + ___ AgCl
13)
___ Ca3P2  ___ Ca + ___ P
14)
___ HCl + ___ F2  ___ HF + ___ Cl2
15)
___ (NH3)2CO3 + ___ CaSO4  ___ CaCO3 + ___ (NH3)2SO4
32
Activity - The Balanced Equation is a Recipe
Introduction
Chemists spend a lot of time and effort making things. Just like a chef spends lots of time and effort making food, a chemist
spends lots of time making chemicals. Chefs use recipes to make food; chemists use recipes to make chemicals.
Purpose
The purpose of this activity is to relate a recipe for brownies to a balanced chemical equation and use this relationship to
determine quantities of ingredients (reactants) and products.
This is my mom’s recipe for one batch of homemade brownies – to be cooked in a 9” x 13” pan.
o 1.5 cups flour (F)
o 2 sticks butter (B)
o 1 tsp salt
o 4 eggs (E)
o 2 tsp vanilla
o 1.67 cups sugar (S) (which
is 1 & 2/3 cups)
o
o
o
4 squares unsweetened
chocolate
¼ tsp baking powder
1 ¾ tsp baking soda
So the recipe goes like this:
2B + 4E + 1.67S + 1.5F (+ more) → 1 batch
1.
If you had 8 eggs (and everything else you need), how many batches of brownies could you make?
2.
Using the 8 eggs, how many sticks of butter would you need to take out of the fridge?
3.
If you had 4.5 cups of flour and wanted to make as many brownies as possible, how many eggs would you need to go
with your 4.5 cups of flour?
4.
You are really hungry. If you had 3 eggs and wanted to make the most brownies that you could, how many cups of sugar
would you need to go with those eggs?
5.
To make 3 batches of brownies, how many cups of flour (and of course all the other necessary ingredients) would you
need? Use the flour to batch ratio to calculate.
6.
2.5 cups of sugar needs how much flour in order to use up all of the sugar?
You see, deciding how much of each ingredient or how many batches you can make is nothing more than manipulating the ratios or proportions
that exist between the ingredients. If you don’t maintain those set proportions, it is likely that your brownies will not taste very good.
333
7.
Write the balanced equation for the synthesis of ammonia (nitrogen trihydride) from nitrogen gas and hydrogen gas.
(Use this balanced equation to solve problems 8 – 11)
8.
Given 7 moles of nitrogen gas, how many moles of hydrogen gas do you need to go with it?
9.
Given 7 moles of nitrogen gas, how many moles of ammonia can you make? Would your answer be the same if you
started with 21 moles of hydrogen?
34
Lesson 6 : Calculating Mole Ratios
Objective: Solve mole-mole Stoichiometry problems given a balanced reaction
Stoichiometry (pronounced as: stow – ik – ee – om′ – etree) is the calculation of quantities in chemical equations. The
question involved with Stoichiometry is “How much…?
For example: based on the balanced chemical equation 3 A2(gs) + 2 B(aq) → 2 A3B(s)
How much:
a) of reactant A is needed to completely react with 5.00 g of reactant B?
b) solid product will form if 0.25 g of reactant A is used?
One of the important concepts in Stoichiometry is that of mole ratios. Mole ratios are the ratios of the
number of moles of each reactant and product to each other.
When we balance chemical equations with the coefficients, we set up the whole number ratios between
reactants, between reactants and each product, and between products.
For example, consider the following balanced chemical equation:
N2 (g)
+ 3 H2 (g)  2 NH3 (g)
Using the coefficients to represent moles of a substance,
Based on the correctly balanced chemical equation, the mole ratios for this reaction are:
ratios between reactants:
1 mole N2
3 moles H2
3 moles H2
1 mole N2
ratios between each reactant and the product:
1 mole N2
2 moles NH3
2 moles NH3
1 mole N2
3 moles H2
2 moles NH3
2 moles NH3
3 moles H2
*****Mole ratios for a given chemical equation are always based on the balanced chemical equation and
NEVER change. ******
For example:
N2 + 3 H2
 2 NH3
1 mole
2 moles
3 moles
0.5 moles
3 moles
6 moles
9 moles
1.5 moles
2 moles
4 moles
6 moles
1.0 moles
Remember: the balanced chemical equation represents the SMALLEST whole number ratios of reactants
and products.
Each of the mole ratios from the original balanced equation may be used as a conversion factor to change
from moles of one substance to moles of another substance in the same chemical equation.
You MUST use the ORIGINAL balanced chemical equation with the correct coefficients as your reference for
mole ratios.
35
Using the reaction of nitrogen gas and hydrogen gas to form ammonia gas;
N2 (g)
+ 3 H2 (g)  2 NH3 (g)
Suppose you had only 0.25 mole N2(g) available. How much hydrogen gas would you need to react
completely with the available nitrogen gas? How much ammonia would be formed?
(1) Start with the amount you are given in the problem
(2) multiply the starting amount by the appropriate mole ratio to convert from the unit with which you
start to the substance desired
(3) cancel units
(4) do the arithmetic and write the answer with appropriate significant figures, correct unit, and correct
chemical formula
0.25 mole N2 x 3 moles H2
1 mole N2
= 0.75 moles H2 are needed
0.25 mole N2 x 2 moles NH3 = 0.50 mole NH3 would be formed
1 mole N2
NOTE: In these calculations, the numbers in the mole ratios do NOT have significant figures. USE THE
NUMERICAL VALUES AS PRINTED IN THE QUESTION AS THE BASIS FOR YOUR SIGNIFICANT FIGURES. In this
calculation, the given value of
0.25 mole N2(g) has 2 significant figures, therefore the answers must each have 2 sig. figs.
More examples:
Consider the reaction between sodium metal and chlorine gas to form solid sodium chloride.
2 Na(s) + Cl2(g) → 2 NaCl(s)
If 25.0 moles of sodium chloride are formed, how many moles of sodium metal were used? How many
moles of chlorine gas were used?
25.0 moles NaCl x
2 moles Na
= 25.0 moles Na were used
2 moles NaCl
25.0 moles NaCl
1 mole Cl2
= 12.5 moles Cl2 were used
2 moles NaCl
x
36
Classwork 5-6:Determine the molar amount using mole to mole ratios for the following reactions
C3H8 + 5O2 → 3CO2 + 4H2O
1. If 12 moles of C3H8 react completely, how many moles of H2O are formed in the reaction above?
C3H8 + 5O2 → 3CO2 + 4H2O
2. If 20 moles of CO2 are formed, how many moles of O2 reacted in the reaction above?
C3H8 + 5O2 → 3CO2 + 4H2O
3. If 8 moles of O2 react completely, how many moles of H2O are formed in the reaction above?
N2 + 3H2 → 2NH3
4. If 2.5 moles of N2 react completely, how many moles of NH3 are formed in the reaction above?
5.
6.
7.
37
8. Use the balanced reaction below and the relationship that “1 mole of a compound = the gram formula mass of that
compound” to answer the questions below.
2H2O  2H2 + O2
a. Calculate the gram formula masses for the three substances seen in the reaction abov
b. How many moles are present in 54 grams of H2O? (Remember: one mole of H2O is ALWAYS equal to 18 grams)
c. What is the ratio of H2O to H2 moles according to the balanced reaction above?
d. Using the reaction above (remember, it is just like a recipe!), how many moles of H 2 would be produced if 4 moles of
H2O are used?
e. How many grams of H2 are present in 4 moles of H2?
f. What is the ratio of H2 to O2 in the reaction above?
g. If you have 2.5 moles of H2O, how many moles of O2 will be produced?
h. What is the ratio of H2O to O2 in the reaction above?
i. If you produce 0.25 moles of O2, how many moles of H2O did you react?
38
Lesson 7: Percent Composition
Objective: To apply a formula to calculate % composition
A common practice in the chemical laboratory is to determine the percent composition of each element in a compound
and use that information to determine the formula that shows the simplest whole-number ratio of elements present in
the compound, the empirical formula of the compound.
Percent composition is defined as
Calculating Percent Composition :Formula located on Table T
EXAMPLE: What is the percent composition of Calcium in CaCl2

Step 1: Calculate the GFM for the compound.
40.1 g
70.9 g
= 111.0 g/mol

Step 2: Use the formula to find the % composition of each element or “part” in our compound (to the nearest
tenth of a %).
%Ca in CaCl2 =40.08 g = 36.11 %
111.0 g
39
*= Honors Chem Only
*% Composition of Hydrates
Water, the most common chemical on earth, can be found in the atmosphere as water vapor. Some chemicals, when
exposed to water in the atmosphere, will reversibly either adsorb it onto their surface or include it in their structure
forming a complex in which water generally bonds with the cation in ionic substances. The water present in the latter
case is called water of hydration or water of crystallization. Common examples of minerals that exist as hydrates are
gypsum (CaSO4•2H2O), Borax (Na3B4O7•10H2O) and Epsom salts (MgSO4•7H2O). Hydrates generally contain water in
specific amounts. The hydrate’s formula are represented using the formula of the anhydrous salt (non-water)
component of the complex followed by a dot then the water (H2O) preceded by a number corresponding to the ratio of
H2O moles per mole of the anhydrous component present. They are typically named by stating the name of the
anhydrous component followed by the Greek prefix specifying the number of moles of water present then the word
hydrate (example: MgSO4•7H2O: magnesium sulfate heptahydrate).
Hydrate- Ionic solids with water trapped in the crystal
lattice. It is written like this:
Ionic Compound’s Formula • n H2O
(n) is a whole number
Notice how WATER molecules are BUILT INTO the chemical formula
Anhydrous salt - A crystal with NO WATER trapped inside its lattice
Calculate % Composition of water in a Hydrate
Step 1: Calculate GFM of the HYDRATE
Step 2: Plug into % composition formula from Table T
% composition by mass = mass water X 100
mass of whole hydrate
Example: What is the percentage by mass of water in sodium carbonate crystals (Na2CO310H2O)?
Step 1- Calculate GFM of Hydrate
=part
=whole
Step 2- Plug values into Formula
% H2O by mass =
****Pay attention because in some problems you will have to calculate the amount of water lost by subtracting the
anhydrate mass from the hydrate mass.
40
Honors Chem Only
Water in a Hydrate using Lab data:
RESULTS:
Data
Remember to record masses to two decimal places:
1. Mass of Crucible
2. Mass of Crucible + Hydrate before heating
3. Mass of Crucible + Anydrate after heating
0.50
_____________g
___________g
__________ g
= 10.0%
5.00
41
Percent of Sugar in Gum
Purpose: Calculate the percent composition by mass of an ingredient in a
commercial food product.
Background: All commercial food producers must list all ingredients found in their food products on the
retail packaging for the foods. The Wrigley Company provides the following nutritional facts for Juicy Fruit
Gum.
Serving size
1 stick (2.7 g)
Calories
10
Total Fat
0g
Sodium
0g
Total carbohydrate
2g
Sugars
2g
Protein
0g
Chewing gum is composed mostly of sugar. The most convenient way to remove sugar from gum
is to chew it. During chewing the sugar is dissolved by saliva and is swallowed, leaving an insoluble gum
base. This laboratory exercise will establish the percent composition of sugar found in chewing gum.
All percentages are calculated the same way. Percent = Part/whole x 100. For example, if a
piece of gum had a mass of 5.0 grams, and the sugar in the gum had a mass of 3.0 grams, the percent of
the gum that was sugar would be calculated by: 3.0/5.0 x 100 = 60 %.
Procedure:
Mass of Gum +
Wrapper (g)
Mass of Wrapper
(g)
Calculated Mass
of Gum alone (g)
1.
2.
3.
4.
Mass of Chewed
Gum + Wrapper
(g)
Calculated Mass
of Chewed Gum
alone (g)
Calculated Mass
of Sugar
Removed after
chewing(g)
Obtain a piece of gum from your teacher. DO NOT OPEN IT YET.
Mass the gum in the wrapper, record the mass in the data table above.
Open the gum, begin chewing. Mass the empty wrapper, record in the data table, and save it.
After chewing until the flavor disappears (or 10 minutes, whichever comes first), place the chewed gum
on your empty wrapper. (While you are chewing the gum, answer the following questions. You
may work together on the questions, but each group member must write the answer on his/her
own paper)
Questions:
 At home, I have a jar of coins. If I have 350 coins in the jar, and 120 of them are quarters, what
percent of the coins are quarters?
42

Determine the percentage of sugar in the gum from the manufacturer’s nutritional statement. *** (Later,
you will use this answer to calculate percent error in your analysis section.)

In this lab, you will mass a piece of gum to find it’s total mass. You will then chew the gum to remove
the sugar. If the sugar is gone after chewing, how will you be able to find the mass of sugar that was in
the gum?
5.
After 10 minutes have passed, re-mass the chewed gum and wrapper together and record the value in
the data table. You may now throw the gum and wrapper away.
6. Complete the calculations that are necessary to fill in your data table.
Analysis of Data:
Calculation for your sample
% sugar in your sample
Group Average (Show work)
1. Determine the percentage of sugar in your sample. Label and show work for your calculations in the
table above. Record your final answer with the correct number of significant figures in the % sugar
column.
2. Determine the average percentage of sugar for your lab group. Show work and record in table above.
3. Determine the percent error between your group average (measured) and the manufacturer’s
nutritional statement that you already calculated in the question ***section (accepted).
4. Explain the percent error that you calculated, including a discussion of possible errors. Remember:
Human or math error is not an acceptable answer.
43
Honors Chem Only
Lab - Percent of Water in Epsom Salt
Goggles must be worn!
Overview:
Epsom salt (aka magnesium sulfate) is a combination of MgSO4 and H2O. Many ionic
compounds incorporate a fixed number of water molecules into their crystal structures. These are
called hydrates. Heat can be used to dehydrate a hydrated salt causing the H2O molecules to
evaporate and produce an anhydrous salt which often will appear different than its hydrate.
When expressing the formula for a hydrate, it is necessary to notate the fixed number of H2O
molecules. A large dot is placed between the formula and the H2O molecules following the formula
for the ionic compound.
For example: CuSO4 • 5 H2O is the formula for copper sulfate. This formula indicates that for
every 1 mole of CuSO4, 5 moles of H2O are present.
Purpose: 1) Dehydrate Epsom salt using chemistry procedures and determine the percent by
mass of water.
Extra Credit 2) Use your math skills to determine the mole ratio, or how many moles of H2O
there are compared to moles of MgSO4 (Epsom salt) 3) Compare your lab determined mole ratio
to the real mole ratio on the box of Epsom salt.
Materials:
Balance
Tongs
Bunsen burner
Ring stand w/ ring
Pipestem triangle
crucible
Epsom salt (MgSO4)
Procedure:
1. Put on your goggles. Secure the iron ring on the ring stand a
couple of inches above the height of the burner. Place
the pipestem triangle on the ring.
2. Place a clean crucible on the set-up. Light
the burner and heat for a couple of minutes to make certain
container is thoroughly dry. Turn off burner and cool the
container for several minutes until it is comfortable to touch.
Record the mass of the dry container.
3. Add about 5 grams of Epsom salt to the container. Record
the mass of the container and Epsom salt.
4. Place the container back on the set-up and heat gently with hot flame until the water has been
released from the hydrate. This will require about 5 minutes. (see illustration)
5. When no more H2O appears to be coming from the hydrate, turn off the burner and cool for
several minutes until container is comfortable to the touch. Record the mass of the container and
salt.
6. If time allows, reheat the container with salt, cool and find the mass again. If the two final
masses agree, you can be confident that you have indeed released all of the H 2O. If not, continue
the heating and cooling as directed in this lab to make sure all the water is released.
7. Put your excess Epsom salt into the trash can and carefully clean out your crucible. All other
lab materials need to be ready for the next class.
44
Data / Observations:
Mass of Epsom Salt before and after heating
1. Mass of a clean, dry, empty container
g
2. Mass of container & hydrated salt
g
(before heating)
3a. Mass of container & anhydrous salt
g
(after heating 1st time)
3b. Mass of container & anhydrous salt (2nd time)
*repeat until no change in mass
g
3c. Mass of container & anhydrous salt (3rd time)
*repeat until no change in mass
g
Calculations/ Analysis (must show work):
1. a. Determine mass (g) of H2O lost from your salt (how much did you “cook out?”): Using your
data table, determine the numbers to subtract in order to figure out mass (g) H2O. Use this number
to calculate the percent by mass of water in the hydrate-Use the formula on Reference Table T.
b. Determine mass (g) of the anhydrous MgSO4 used in this lab (just the anhydrous salt- not
the molar mass). Using your data table, calculate the mass of the anhydrous MgSO4 used.
Extra Credit (must show work):
2. a. Calculate how many moles of H2O were dehydrated from your salt.
Convert your grams of H2O (found in 1a) to moles of H2O.
b. Calculate how many moles of MgSO4 salt. Convert your grams of
MgSO4 (found in 1b) to moles of MgSO4.
3. a. Determine the mole ratio of H2O molecules to MgSO4 particles. Divide the moles of H2O
(calculated in 2a) by the moles of MgSO4 (calculated in 2b) to determine the ratio of moles of
H2O to moles of salt (moles H2O / moles MgSO4) . Round your answer to the nearest small
whole number.
b. Write the correct formula for your hydrate using the ratios
your answer from #3 to fill in the blank).
1MgSO4•_____H2O (use
Conclusion/Discussion:
1. Copper Nitrate is a hydrate with the following formula: Cu(NO3)2 • 3 H2O. What is the ratio
between moles of copper nitrate and moles of water in this hydrate?
2. What is the percent by mass of water in the hydrate in Cu(NO3)2 • 3 H2O? Show All Work!!!
45
Classwork 5-7: Show all calculations - Round percentages to whole number
1. What is the percentage by mass of carbon in CO2?
2. What is the percent by mass of nitrogen in NH4NO3?
3. What is the percent by mass of oxygen in magnesium oxide(MgO)?
4. What is the percent by mass of water in BaCl 22H2O?
Question 4 & 5 - Honors Chem
Only
5. A 10.40 gram sample of hydrated crystal is heated to a constant mass of 8.72 grams. This
means all of the water has been driven out by the heat.
a) Calculate the mass of water that was driven out:
b) Calculate the %mass of water in the hydrate.
46
Chemical Calculations Review
47
48
49
50
Unit Review:
Moles and Chemical Reactions
Place a checkmark next to each item that you can do! If a sample problem is given, complete it as evidence.
Place a checkmark next to each item that you can do! If a sample problem is given, complete it as evidence.
_____1. I can still do
everything from Unit 1.
_____2. I can still do
everything from Unit 2.
_____3. I can still do
everything from Unit 3.
_____4. I can still do
everything from Unit 4.
_____5. I can still do
everything from Unit 5.
a. O2
______6. I can calculate the
gram formula mass of a
compound or substance
b. Na3PO4
a. 64g O2
______7. I can calculate the
number of moles of a substance b. 567 g Na3PO4
when given the mass
a. 7 moles O2
______8. I can calculate the of
the mass a substance when
given number of moles
b. 0.6 moles Na3PO4
51
Definitions:
empirical formula
_____9. I can define empirical
formula, molecular formula,
and hydrate.
molecular formula
What is the molecular formula of a compound that has the empirical
formula of CH and a molar mass of 78 g/mol.
_____10. Given the empirical
formula and the molar mass, I
can determine the molecular
formula of a compound.
Using the symbols shown below, complete the equation below to
illustrate conservation of mass.
= Al
_____11. I can use particle
diagrams to show conservation
of mass in a chemical equation.
_____12. I can balance a
chemical equation showing
conservation of mass using the
lowest whole number
coefficients.
= Br
Br
2Al
+
3Br2
----->
2AlBr3
Balance the following chemical equation using the lowest whole number
coefficients.
_____Al2(SO4)3 + _____Ca(OH)2 -----> _____Al(OH)3 +
_____CaSO4
Use the law of conservation of mass to predict the missing product.
_____13. Given a partially
balanced equation, I can
predict the missing reactant or
product.
2NH4Cl + CaO -----> 2NH3 + ______________ + CaCl2
52
_____14. Given a list of
chemical reactions, I can
classify them as being a
synthesis reaction,
decomposition reaction, single
replacement reaction, or
double replacement reaction.
Classify the following reactions as synthesis, decomposition, single
replacement, or double replacement.
Given the following balanced equation, state the mole ratios between the
requested substances.
C3H8(g) + 5O2(g) -----> 3CO2(g) + 4H2O(l)
_____15. Given a balanced
equation, I can state the mole
ratios between any of the
reactants and/or products.
The mole ratio between C3H8 and O2 is _______C3H8:_______O2.
The mole ratio between C3H8 and CO2 is _______C3H8:_______CO2.
The mole ratio between C3H8 and H2O is _______C3H8:_______H2O.
The mole ratio between CO2 and O2 is _______CO2:_______O2.
The mole ratio between H2O and CO2 is _______H2O:_______CO2.
Definition:
stoichiometry
_____16. I can define
stoichiometry.
Using the equation from question #20, determine how many moles of O2
are needed to completely react with 7.0 moles of C3H8.
_____17. Given the number of
moles of one of the reactants or
products, I can determine the
number of moles of another
reactant or product that is
needed to completely use up
the given reactant/product.
Using the equation from question #20, determine how many moles of
CO2 are produced when 7.0 moles of C3H8 completely react.
53
Unit 5 Practice Test
1) How many oxygen atoms are represented in the formula Al2(CO3)3?
a) 3
b) 9
c) 10
d) 6
2) What is the total number of atoms present in 1 mole of Ca3(PO4)2?
a) 8
b) 5
c) 10
d) 13
3) What is the total mass of iron in 1.0 mole of Fe2O3?
a) 72 g
b) 112 g
c) 56 g
d) 160 g
4) What is the gram formula mass of Li2SO4?
a) 206 g
b) 55 g
c) 110 g
d) 54 g
5) What is the percent composition by mass of sulfur in H2SO4? [formula mass = 98]
a) 98%
b) 16%
c) 65%
d) 33%
6) A hydrate is a compound with water molecules incorporated into its crystal structure. In an
experiment to find the percent by mass of water in a hydrated compound, the following data
were recorded:
Mass of test tube + hydrate crystals before heating
Mass of test tube
Mass of test tube + anhydrate crystals after heating
25.3 grams
21.3 grams
22.3 grams
What is the percent by mass of water in the hydrate?
a) 75%
b) 50%
c) 8.0%
d) 95%
7) Which of the following statements explains why mass is lost when a student heats a sample of
BaCl22H2O crystals?
a) water is given off as a gas
b) the crystals sublime
c) chlorine is given off as a gas
d) the crystals fuse (melt)
8) When the equation H2S + O2  H2O + SO2 is completely balanced using the smallest whole
numbers, the sum of all the coefficients is
a) 9
b) 11
c) 7
54
d) 5
9) What is the percent by mass of water in the hydrate Na2CO3●10H2O [formula mass = 286]?
a) 26.1%
b) 62.9%
c) 6.89%
d) 214.5%
10) How many grams are there in 2.5 moles of C2H5OH?
a) 18
b) 115
c) 46
d) 0.05
11) Which quantity is equivalent to 146 grams of NaCl?
a) 1.0 mole
b) 2.5 moles
c) 2.0 moles
d) 1.5 moles
12) When the equation H2 + Fe3O4 --> Fe + H2O is completely balanced using the smallest
whole numbers, the coefficient of Fe would be
a) 1
b) 2
c) 3
d) 4
13) When the equation NaBr + H3PO4  Na3PO4 + HBr is balanced using the smallest whole
numbers, the sum of the coefficients will be
a) 6
14) Given the equation:
b) 8
c) 5
d) 4
2C2H2 + 5O2  4CO2 + 2H2O
How many moles of oxygen are required to react completely with 1.0 mole of C2H2?
a) 2.5
15) Given the reaction:
b) 5.0
c) 10
d) 2.0
2C2H6 + 7O2  4CO2 + 6H2O
What is the ratio of moles of CO2 produced to moles of C2H6 consumed?
a) 7 to 2
b) 2 to 1
c) 1 to 1
d) 3 to 2
16) Given the following balanced equation: N2 + 3 H2  2 NH3
If you have 8 moles of nitrogen, how many moles of NH3 will you produce?
a) 16
b) 18
c) 24
d) 2
17) Given the reaction:
2Na + 2H2O  2NaOH + H2, what is the total number of
moles of hydrogen produced when 4 moles of sodium react completely?
a) 1
b) 2
c) 3
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d) 4
18) What type of reaction best describes the following chemical reaction?
Zn + CuSO4 --> ZnSO4 + Cu
a) single replacement
b) double replacement
c) decomposition
d) synthesis
19) Which chemical equation best represent a decomposition reaction?
a) Cl2 + 2KI  2KCl + I2
c) H2CO3  H2O + CO2
b) 2Al + 3Cl2  2AlCl3
d) KCl + AgNO3  KNO3 + AgCl
20) What is the molecular formula of a compound that has a molecular mass of 54 and an
empirical formula of C2H3?
a) C8H12
b) C6H9
c) C4H6
d) C2H3
21) What is the empirical formula of the compound whose molecular formula is P4O10?
a) P8O20
b) PO2
c) P2O5
d) PO
22) Which of the following is an empirical formula?
a) H2O2
b) H2O
c) C2H2
d) C4H8
23) Which represents both an empirical and molecular formula?
a) P2O5
b) C3H6
c) N2O4
d) C6H12O6
For questions 24 and 25, show all work and express your answer in the appropriate units.
24) Calculate the gram formula mass of ZnSO4.
25) Use your answer from 23 to calculate the percent by mass of zinc in ZnSO4.
26) Balance the following reaction and reduce to the lowest whole number coefficients
_____ H2SO4
+
_____ B(OH)3

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_____ B2(SO4)3
+ _____ H2O
27) Li and KNO3 according to the following equation:
Li + KNO3  LiNO3 + X
Write the formula for the missing product X.
Use the chemical equation below to answer questions 28-30.
___ N2 +
___ H2

___ NH3
27) Balance the equation above using the lowest whole number coefficients.
29) How many moles of H2 are required to produce 6.5 moles of NH3? Show all work and make
sure your answer has the correct number of significant figures and proper units.
30) How many grams of NH3 are produced if 50.0 grams of N2 are consumed? Show all work and
make sure your answer has the correct number of significant figures and proper units.
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