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Marking Period # 3 (Enduring Understanding:___________________)
California State Standard
Conservation of Matter and Stoichiometry
3. The conservation of atoms in chemical reactions leads to the principle of conservation of
matter and the ability to calculate the mass of products and reactants. As a basis for
understanding this concept:
a. Students know how to describe chemical reactions by writing balanced equations.
b. Students know the quantity one mole is set by defining one mole of carbon 12 atoms to
have a mass of exactly 12 grams.
c. Students know one mole equals 6.02 x 1023 particles (atoms or molecules).
d. Students know how to determine the molar mass of a molecule from its chemical
formula and a table of atomic masses and how to convert the mass of a molecular
substance to moles, number of particles, or volume of gas at standard temperature and
pressure.
e. Students know how to calculate the masses of reactants and products in a chemical
reaction from the mass of one of the reactants or products and the relevant atomic
masses.
f. *Students know how to calculate percent yield in a chemical reaction.
g. *Students know how to identify reactions that involve oxidation and reduction and how
to balance oxidation-reduction reactions.
*Not tested on CST
NOTE: DOCUMENTS (1) TO (18) ARE STORED IN MP 3 SUPPLEMENTARY FILE
Day 1 (Essential Question(s): ____________________________________________________________________________________)
Learning Objectives:
Define mole.
Distinguish between types of particles (atoms, formula units/ions, molecules)
Use dimensional analysis to convert between moles and number of particles.
Standard 3b, c
Activity
Warm Up:
1.Inspect and
2.Record mass of one
mole of Pb, Cu, Zn, S,
NaCl, sugar and water
3.Answer Q1and Q2
Avogadro’s No. the
concept of 1 mole
=6.02X1023 Use
dimensional analysis
for Mole Conversion
Check for
Understanding
Objective
Describe and interpret
one mole of atoms,
formula units (f.u.),
molecules as related to
the mass of elements and
compounds
Convert between moles
and particles using
conversion factors:
Materials Required
Samples of
(a)207.2g Pb,(b)63.55g Cu;
(c)65.39g Zn, (d)32.07g S;
(e)58.44g NaCl;
(f)342.30gsugar
(C12H22O11)
(g)18g water(H2O)
7 different jars
Pass around classroom
for observations
Time
15 minutes
(1) Lecture Notes and
Worksheet on Mole to
Particles Conversion
35 minutes
1 /2mole of H2O= ? molecules
1.2X1024 molecules of
H2O=?mole
5 minutes
1mole/6.02X1023particles
6.02X1023particles/1mole
Notes
Discovery Demo:
Q1.What is the
particle/unit for each
substance?
elements(atoms),
ionic cpds (f.u.),
covalent/molecular
cpds (molecules)
Q2. What is the
significance of the
mass of each
substance? (Hint: use
periodic table)
Students write
answers of every
question on board
show each step
Class Discussion of
student work
(Prentice Hall)HW:
P.344
#31;36a,b;37a,b;
38a,b;39a,b
Day 2 (Essential Question(s): ____________________________________________________________________________________)
Learning Objectives: 3a, d


Distinguish and calculate atomic mass, formula mass, molecular mass and molar mass.
Convert mass of elements, ionic, molecular compounds to mole.
Activity
1.Show 1
mole= 12 g of
Carbon
2.Review 1
mole of the
different
substances as
related to
atomic masses
Power points
on conversion
of molar mass
of elements,
ionic
molecular,
compounds to
moles
Check for
understanding
Objective
1.Students know the quantity of
one mole is set by defining one
mole of C- 12 atoms to have a
mass of exactly 12 grams.
2. Students know how to
determine the molar mass of a
molecule from its chemical
formula and a table of atomic
masses
1.Distinguish the difference
between molar mass, formula
mass and molecular mass
2.Calculate molar mass/formula
mass/molecular mass
3.Convert mass of elements,
ionic and molecular compounds
to moles
Materials Required
12g Carbon
And Previous day samples
(13) Diagram: Formula Mass of
H2O
Worksheet (2) + Answers (3) on
Molar Mass Calculation
(4) Lecture figures on mass to
mole conversion
(7) Mass to Mole Power points
(A)
(a)1/2 mole C =? Atoms= ? g
(b) 1.20 g of Carbon = ? mole
(c) 5.5 g of NaCl=? Mole= ? f.u.
Time
10 minutes
Notes
Review mole,
particles and
atomic masses
40 minutes
Work on
Practice
Questions
Worksheet(2)
5 minutes
Students show
work on board
and class
discussion
HM Wk: P.344
#30;#32a,b;
#33,a,b;
#34,a,b;
#35,a,b
Day 3 (Essential Question(s): ____________________________________________________________________________________)
Learning Objective:
Convert mass of elements, ionic compounds and covalent/molecular compounds to moles to number of particles.
Standard 3b, c, d
Activity
Objective
Materials Required
Review mole Convert moles of elements, ionic (8)Mass to Mole Power Points (B)
to mass & mole compounds
and
molecular
to
Particles compounds to mass and particles
Conversions
Convert mass
Summarize Mole Conversion and
(5) Lecture Notes on Mole
to mole to
expand mass to mole to particles
Conversion
particles
conversions
(5A) In-class mole assignment
Time
10 minutes
Notes
5.50g NaCl=?
mole
5.50X1024=? fu
40 minutes
Work on
practice
problems in (5)
& Answers
(5A)
Students show
all work for (5)
on board
6 minutes
Hm Wk: 10:2
Worksheet#2,
3,67,10,11,14,
16,17
Practical Example of Conversion
from mass to fu
Check for
understanding
Discuss results based on student
work
(9) 10:2 Worksheet
Day 4 (Essential Question(s): ____________________________________________________________________________________)
Learning Objectives:
 Convert mass to number of moles to number of particles.
Standard 3b, c, d
Activity
Molar
Quantities Lab
or Hydrate Lab
or Chalk MiniLab
(activity sheets
shown below)
Check for
Understanding:
Class
Discussion of
Lab Results and
Questions
Objective
Practice mass to mole to
particles conversion
Materials Required
(10) Molar Quantities Lab
Strips (similar size)of Pb, Cu, Zn,
and about
5 g of S, NaCl, sugar, H2O,
electronic scales
Time
45 minutes
Notes
Complete Lab
and Answer
Lab questions
10 minutes
Hm Wk: 10:2
Practice
#18 to 22
Lab: Determining the Formula of a Hydrate
Introduction
When certain ionic solids crystallize from aqueous solutions, a definite number of molecules of water
remain attached to the crystal. Ionic solids that contain a definite amount of water are called hydrates or
hydrated salts and the water in the crystal structure is called water of hydration. The water is loosely
bound to the ionic solid so it is possible to dehydrate or remove the water by heating. The solid that
remains after all of the water is removed is said to be anhydrous.
In this experiment, the formula of a hydrate, CuSO4∙xH2O, will be determined. A known mass of hydrated
copper sulfate will be heated to remove all of the water. From the mass of the solid before and after
heating, the number of moles of water of hydration, x will be calculated.
Pre-Lab Questions:
(Answer on a separate piece of paper).
Washing soda is a hydrated compound whose formula can be written Na2CO3∙xH2O, where x is the
number of moles of H2O per mole of Na2CO3. When a 2.123 g Na2CO3∙xH2O was heated at 130oC, all of
the water of hydration was lost, leaving 0.787 g Na2CO3.
(1) Calculate the mass of water lost.
(2) Convert the mass of water lost to moles.
(3) Convert 0.787 g Na2CO3 to moles.
(4) What is the ratio of moles of H2O to moles of Na2CO3?
Materials:
about 2 g of copper sulfate
clay triangle
crucible with cover
crucible tongs
ring stand
iron ring
balance Bunsen burner
Procedure:
1. Prepare the heating set-up.
2. Heat a clean crucible strongly for a minute. Allow it to cool.
3. Mass the crucible.
4. Place about 2 g of copper sulfate into the crucible and immediately mass the crucible containing
the hydrate.
5. GENTLY heat the crucible. If temperature is too high, the hydrated crystals may spatter. After 2
minutes, increase the intensity of the flame slightly. Continue heating until the blue color
completely disappears.
6. Allow the crucible to cool and mass it.
7. To make sure that all of the water is removed, repeat steps 5-6 until the mass of the crucible and
its contents stays the same.
8. Observe the contents of the crucible. Add a few drops of water.
Data and Observations:
Mass of crucible (g)
Mass of crucible and hydrated copper sulfate (g)
Mass of crucible and anhydrous copper sulfate (after1st heating) (g)
Mass of crucible and anhydrous copper sulfate (after 2nd heating) (g)
Mass of hydrated copper sulfate (g)
Mass of anhydrous copper sulfate (g)
Mass of water lost (g)
Observations:
Calculations:
(1) Calculate the moles of water lost.
(2) Calculate the moles of anhydrous copper sulfate (CuSO4).
(3) Calculate the moles of water lost per mole of anhydrous copper sulfate.
Analysis and Conclusion:
1. Why must the crucible be cooled before massing?
2. What happened when you added water to the anhydrous solid? What does this indicate?
3. The correct formula for hydrated copper sulfate is CuSO4∙5H2O. Did you get the same value of x?
If not what could be some possible sources of errors?
4. What is the chemical name of CuSO4∙5H2O?
Mole Mini-Lab
How many moles of chalk is your NAME worth?
1. Get a piece of chalk and mass it.
Initial mass of chalk = __________
2. Use the chalk to write your whole name on the pavement.
3. Mass your chalk again.
Final mass of chalk = ___________
4. Do the following calculations:
(a) How many grams of chalk did you use?
(b) Chalk is calcium carbonate, CaCO3. What is the molar mass of CaCO3?
(c) Convert the mass of CaCO3 to moles.
5. So how many moles of chalk is your name worth?
__________________________
For extra credit, calculate the number of ions of Ca2+ and CO32- that you used to write your name.
DAY 5
VETERAN’S DAY HOLIDAY
Day 6 (Essential Question(s): ____________________________________________________________________________________)
Learning Objectives:
Review conversion of mass to number of moles to number of particles.
Define molar volume.
Convert molar volume to number of moles to mass of a gas at STP.
Standard 3b, c, d
Activity
1.Review on
Mole
Conversions
Objective
Review conversion from Mass to
Mole to Particles
2.Molar
Volume
Conversion
Define STP
Define Molar Volume as 22.4L at
STP
Conversion of Volume to Mole to
Mass
Check for
understanding
Students Recreate Mole Map to
enhance Mole Conversion
Materials Required
(11) Mole Map
(6) Mole Conversion
Worksheet
(6A) Answers for (6)
(14) Volume to Mole and
Mole Map for Review
(6) Mole Conversion
Worksheet
(9) 10:2 Practice
Time
20 minutes
Notes
Use Mole Map
to guide
conversion
30 minutes
Do (6) in class
and HW: (9)
10:2 Practice
Problems
Worksheet
#23 to #30
Draw Mole
Map
Check for X or
/ for mole
conversion
based on Mole
Map
5 minutes
Day 7 (Essential Question(s): ____________________________________________________________________________________)
Unit Test Review (Version 1)
Standard 3 a,c,d
Activity
Review concepts 3a,c,d
Objective
Prepare students to
take tests: recognize
what is given and what
is asked for the
different mole
conversions
Practice
Skill
Test
Materials Required
(15) and (16) Review
Questions
Similar to questions in
STAR/SAT tests
Time
56 minutes
Notes
Students write answers
on board and show
work
Class discussion of all
problems
Taking
Day 7 Unit Test Review (Version 2)
A PowerPoint is attached. Review may be done using a Cooperative Learning strategy called
Numbered-Heads-Together. Students working in groups of 3 figure out answer to each
question together. Group member whose number is called writes the answer on white board
for teacher to check.
A Review Sheet is also attached.
Slide 1
Slide 6
Test Review
How many particles are there in
0.25 moles of a substance?
Mole Concept
1.5 x 1023 particles
Slide 2
Slide 7
Which of these is true?
A. Different substances with the same number
of moles have the same mass.
B. One mole of any substance contains the
same number of particles.
C. Mole is a unit of energy.
D. One mole of any substance contains
6.02 x 1021 particles.
What is the mass of 6.02 x 1023
molecules of CO2?
44 g
B
Slide 3
Slide 8
What is the molecular mass of
C4H10?
Which is equal to 45 g of H2O?
A. 1 mole
B. 1.5 moles
C. 2.0 moles
D. 2.5 moles
58 amu
D
Slide 4
Slide 9
What is the molar mass of HNO3?
What is the molar mass of CaCl2?
63 g
110 g
Slide 5
Slide 10
How many moles are there in
60.0 g of carbon?
What is the mass of 0.75 moles
of CaCl2?
83 g
5.00 moles
Review: Mole Concept
This test will evaluate how well you can do the following:
1. Calculate the formula, molecular or molar mass of a substance.
Sample Questions:
A. What is the molecular mass of hydrogen peroxide (H2O2)?
B. What is the molar mass of calcium chloride (CaCl2)?
2. Perform mass-mole, mole mass, mass-mole-number of particles conversion.
Sample Questions:
A. How many moles are in 75.0 g of water (H20)?
B. What is the mass of 1.5 moles of sugar (C12H22O11)?
C. How many formula units (particles) of sodium chloride (NaCl) are in 100.0 g of this
substance?
Day 8 Unit Test
A Unit Test is attached below. Tweak it as you see fit.
Unit Test: Mole
Part 1. Multiple Choice. Circle the letter of the correct answer. For questions marked with an
asterisk (*), show your work.
1) *Which of these substances has a molar mass of 64.0 g?
A. O2
B. CH3OH
C. SO2
D. CaCl2
2)
What is the molar mass of CaCl2?
A. 110. g
B. 90.0
C. 75.0 g
D. 70.0 g
3) *How many moles are contained in 9.03 x 1023 molecules of oxygen gas (O2)?
A. 1.00 mole
B. 1.50 moles
C. 2.00 moles
D. 2.50 moles
4) *How many moles are contained in 45.0 g of H2O?
A. 1.00 mole
B. 1.50 moles
C. 2.00 moles
D. 2.50 moles
5) *What is the mass of 0.750 moles of potassium chloride, KCl?
A. 149 g
B. 74.5 g
C. 55.9 g
D. 37.3 g
6) Standard temperature and pressure (STP) is
A. 0oC and 2 atm
B. 100oC and 1 atm
C. 0oC and 1 atm
D. 100oC and 2 atm
7) What is volume of a gas at STP?
A. 1.0 L
B. 2.4 L
C. 22.4L
D. 44.8 L
8) *How many moles of a gas occupy a volume of 33.6 L at STP?
A. 1.00 mole
B. 1.50 mole
C. 2.00 mole
D. 3.00 mole
9) Which is true about one mole of calcium nitrate, Ca(NO3)2?
A. It has a mass of 116 g.
B. It has a mass of 164 amu.
C. It contains 6 oxygen atoms.
D. It contains 1.204 x 1024 nitrate (NO3-) ions.
10) 9.03 x 1023 atoms of silver are placed on a balance. The balance should read
A. 53.96 g
B. 107.87g
C. 161.81g
D. 215.74g
Part 2. Free Response. Answer the questions as comprehensively as you can. Make sure that
calculations have correct units and correct number of significant digits.
A student was tasked to determine the number of moles of water (n) in one mole of MgCl2·nH2O. She
placed a small sample of MgCl2·nH2O in a dry crucible and heated it several times until all of the water
has evaporated. From the mass before and after heating, she was able to determine the mass and the
number of moles of water in the sample. The chart below shows the data she gathered.
Mass of empty container
Initial mass of sample and container
Mass of sample and container after first heating
Mass of sample and container after second heating
Mass of sample and container after third heating
22.347 g
25.825 g
23.982 g
23.976 g
23.976 g
(1) Explain why the student can correctly conclude that the hydrate was heated a sufficient number of
times in the experiment.
_____________________________________________________________________________________
_____________________________________________________________________________________
(2) Use the data above to
(i) calculate the mass of the water that was lost upon heating
(ii) calculate the number of moles of water lost when the sample was heated
(iii) calculate the mass of MgCl2 that remain in the crucible.
(iv) calculate the mole of MgCl2 that remain in the crucible.
(v) How many moles of water are lost per mole of MgCl2? What is the formula of the hydrate?
Day 9 (Essential Question(s): ____________________________________________________________________________________)
Learning Objectives:
 Recognize signs/evidences of a chemical reaction.
 Recognize the reactants and products in a reaction.
 Identify different types of chemical reactions.
Standard 3a
Activity
Warm Up
Direct Instruction and
Guided Practice Evidences of Chemical
Change, Reactants and
Products, Types of
Reaction
Check for
Understanding/Exit
Ticket
Objectives
Review the difference
between physical and
chemical change.
Students identify evidences
of a chemical change,
identify reactants and
products and classify
reactions.
Materials Required
Time
5 minutes
Notes
Physical or Chemical
Change?
1. melting of ice
2. rotting of food
3. burning of gasoline
4. evaporation of water
Chemical Reactions and
Stoichiometry PowerPointsee Supplementary
Materials folder
40 minutes
The sample reactions
could be shown on
PowerPoint or through a
demo. Reactions include:
(1) Mg + HCl
(2) burning of paper
(3) Pb(NO3)2 + KI
Demo: Mg, HCl solution,
paper, matches, Pb(NO3)2
and KI solution, beaker,
test tubes, droppers
NaHCO3 + CH3COOH →
CH3COONa + H2CO3
10 minutes
NaHCO3 + CH3COOH →
CH3COONa + H2O + CO2
The equations above show
the reaction between
baking soda and vinegar.
1) What evidences of a
chemical reaction can be
observed as the reaction is
occurring?
2) What are the reactants
of the reaction? What are
the products?
3) What type of chemical
reaction is it?
Homework
Practice Worksheet
Questions 1-3
See below -Chemical
Reactions and
Stoichiometry Practice
Worksheet
The Practice Worksheet is
given at the beginning of
the unit. Certain
questions are assigned
per day for students to
practice on.
PowerPoint Slide Master
Slide 1
Slide 5
Signs of Chemical Reactions
Chemical Reactions and
Stoichiometry
Formation of a
precipitate
Precipitate –
insoluble solid
formed from the
reaction between 2
aqueous solutions
Slide 2
Slide 6
Objectives:
 Recognize signs of chemical reactions.
Chemical Reaction and Equation
Chemical reaction – a change that forms new substances
 Recognize the reactants and products in a
Reactants – starting substances
Products – new substances formed
reaction.
 Identify different types of chemical reactions.
Chemical Equation
- Shorthand way of describing chemical reactions
Example:
2H2
+
(Reactants)
Slide 3
Signs of Chemical Reactions
O2
→
2H2O
(Product)
Slide 7
AgNO3 + NaCl → AgCl + NaNO3
What are the reactants in the above reaction?
Evolution of a gas
What are the products?
Slide 4
Signs of Chemical Reactions
Slide 8
Types of Chemical Reactions
1.
Change in
intensive
properties like
color, odor,
density
2 or more reactants
→
Example: H2 + O2 → H2O
Types of Chemical Reactions
3. Single Displacement
2 reactants
1 product
2. Decomposition
1 reactant →
2 or more products
Example: H2O2 → H2O + O2
Release or
absorption of
energy
Slide 9
Combination or Synthesis
Slide 11
Types of Chemical Reactions
Classify each reaction:
→
2 products
(active element and compound)
Example: Mg + HCl → MgCl2 + H2
4. Double Displacement
2 reactants
→
2 products
1. NaCl + AgNO3 → AgCl + NaNO3
2. Na + H2O → NaOH + H2
3. C2H6 + O2 → CO2 + H2O
4. Mg + O2 → MgO
(2 aqueous solutions)
5. Na2CO3 → Na2O + CO2
Example: KI + Pb(NO3)2 → KNO3 + PbI2
6. KOH + HCl → KCl + H2O
Slide 10
Types of Chemical Reactions
5. Combustion
fuel + oxygen
→
water + carbon dioxide
Example: CH4 + O2 → CO2 + H2O
Chemical Reactions and Stoichiometry
Practice Worksheet
1. Describe the different signs of chemical reactions:
A. _________________________________
B. _________________________________
C. _________________________________
D. _________________________________
2. Classify the following reactions as DECOMPOSITION, COMBINATION, SINGLE
DISPLACEMENT, DOUBLE DISPLACEMENT and COMBUSTION.
(a) CaCO3 → CaO + CO2
____________________________________
(b) BaCl2 + Na2SO4 → BaSO4 + 2NaCl
____________________________________
(c) 3HNO3 + Al(OH)3 → 3H2O + Al(NO3)3
____________________________________
(d) 2C2H2 + 5O2 → 4CO2 + 2H2O
____________________________________
(e) Na2O + H2O → 2NaOH
____________________________________
(f) Mg + 2HCl → MgCl2 + H2
____________________________________
3. Name the type of reaction described below.
___________________ (a) a complex compound breaks down into simpler compounds or into its
constituent elements
___________________ (b) two or more elements or simpler compounds react to form a single more
complex compound
____________________ (c) a more active element displaces a less active one from its compound
____________________ (d) reaction between two solutions of ionic compounds
____________________ (e) reaction that requires oxygen as a reactant and produces carbon dioxide and
water
4. Write a chemical equation for each chemical reaction described below.
A.) Magnesium reacts with hydrochloric acid to produce magnesium chloride and hydrogen gas (H2).
What are the reactants? ________________________________________________________
What are the products? ________________________________________________________
Chemical Equation: ___________________________________________________________
B.) Iron reacts with oxygen (O2) in air and forms iron(III) oxide.
What are the reactants? ________________________________________________________
What are the products? _________________________________________________________
Chemical Equation: ____________________________________________________________
C.) Two clear, colorless solutions of potassium iodide and lead(II) nitrate react with each other and
produce potassium nitrate and lead(II) iodide, a yellow precipitate.
What are the reactants? _______________________________________________________
What are the products? _______________________________________________________
Chemical Equation: __________________________________________________________
5. What type of reaction will most likely occur to the given reactant(s)? Complete the equation by
predicting the products of the reaction.
A) Fe(s) + HCl(aq) →
B) AgNO3(aq) + NaCl(aq) →
C) Mg(s) + CuSO4(aq) →
D) Ag2O →
E) KOH(aq) + BaCl(aq) →
F) C4H10(g) + O2 (g) →
G) Na(s) + O2(g) →
6. Complete the paragraph.
According to the Law of __________________________________, mass remains the same
before and after a chemical reaction. This is because atoms are not ___________________________ nor
_______________________ during a chemical reaction. The number and kind of atoms do not change.
This is shown in a balanced chemical equation where the same number of atoms is written on each side of
the equation. To balance an equation, ________________________ are written before the formula of the
reactants and products.
7. Balance these chemical equations:
A.)
Ca
B.)
H2O2
C.)
N2
D.)
E.)
+
O2
→
→
H2
H2
→
Cu2O
+
C
CO2 +
H2O →
+
CaO
+
O2
NH3
→
Cu
+
CO2
C6H12O6
+
O2
Solve the following stoichiometric problems. Show your work.
8. Mole-Mole Problems
In the chemical reaction,
Mg
+
2HCl →
MgCl2
+
A. how many moles of magnesium are needed to produce 3.00 moles of hydrogen gas?
H2
B. how many moles of magnesium chloride can be produced from 4.00 moles of hydrogen chloride?
In the reaction,
4Fe
+
3O2
→
2Fe2O3
C. how many moles of oxygen are needed to react with 2.00 moles of iron?
D. how many moles of iron and oxygen are needed to produce 6.00 moles of iron (III) oxide?
9. Mole-Mass Problems
In the chemical reaction,
Mg
+
2HCl →
MgCl2
+
A. how many moles of magnesium are needed to produce 4.00g of hydrogen gas?
H2
B. how many grams of magnesium chloride can be produced from 2.00 moles of magnesium?
In the reaction,
4Fe
+
3O2
→
2Fe2O3
C. how many grams of iron are needed to form 4.00 moles of iron (III) oxide?
D. how many moles of oxygen are needed to completely react with 112 g of Fe?
10. Mass-Mass Problems
In the chemical reaction, C6H12O6(g) + 6O2(g)
→
6CO2(g)
+
6H2O(g)
A. how many grams of carbon dioxide can be produced from the burning of 180.0 g of glucose
(C6H12O6)?
B. what is the mass of oxygen needed to produce 54.0 g of water?
In the chemical reaction,
2Mg +
CO2 →
2MgO +
C. how many grams of carbon dioxide are needed to produce 36.0 g of C?
C
D. what is the mass of magnesium oxide that can be produced from 36.0 g of magnesium?
11. Molar Volume
In the chemical reaction,
Mg(s) +
2HCl(aq)
→
MgCl2(aq)
+
A) how many liters of hydrogen gas at STP is produced from 2.50 moles of magnesium?
H2(g)
B) how many liters of hydrogen gas at STP is produced from 18.0 g of magnesium?
In the chemical reaction, C6H12O6(g) + 6O2(g)
→ 6CO2(g)
+
6H2O(g)
C) How many liters of oxygen are needed to produce 12.0 moles of carbon dioxide at STP?
D) How many liters of carbon dioxide are produced from 16.0 L of oxygen at STP?
12. Limiting and Excess Reactants
A)
2 slices of bread
+
3 slices of ham
For the burger “reaction”, complete the table below:
Number of
Number of Ham Number of
Bread Slices
Slices
Sandwiches
10
20
20
24
6
6
20
12
10
12
→
Name of
Limiting
“Reactant”
2 sandwiches
Name of Excess
“Reactant”
Excess amount
ham
20 slices
bread
10 slices
B) Hydrogen gas reacts with oxygen gas to form water vapor according to the reaction below:
2H2(g) +
O2(g) →
2H2O(g)
For this reaction, complete the table below:
Amount of
Hydrogen
Amount of
Oxygen
10 moles
8 moles
20 moles
6 moles
4g
40 g
10 g
64 g
2L
2L
12 L
4L
Amount of
Water Vapor
Name of
Limiting
Reactant
Name of Excess
Reactant
Excess Amount
13. Theoretical and Percent Yield
A. In the chemical reaction,
2Mg +
CO2 →
2MgO +
C
if only 58 g of MgO is actually produced from 36.0 g of Mg, what is the percent yield of the reaction?
B. What is the percent yield of the reaction shown below if 11.0 g of hydrogen reacts completely with
nitrogen to form 40.8 g of ammonia?
N2
+
3H2
→
2NH3
Day 10 (Essential Question(s): ____________________________________________________________________________________)
Learning Objectives:
 Write chemical equations from word equations.
Standard 3a
Activity
Warm Up
Objectives
Review identifying reactants and
products and classifying reactions.
Materials Required
Time
5 minutes
Notes
Al + O2 → Al2O3
1. What is/are the
reactant(s) of the
reaction shown
above?
2. What type of
reaction is it?
Explain.
Direction Instruction:
Predicting Products
Students view the reaction between
sodium and chlorine. They identify the
reactants and product of the reaction.
They will then be guided in writing the
chemical equation for the reaction.
Youtube video of
reaction between
sodium and chlorine –
see link below
10 minutes
Guided Practice
Students look at other chemical
reactions. Working in groups, they
practice identifying reactants and
products and writing chemical equations.
Chemical Reactions
and Stoichiometry
PowerPoint- see
Supplementary
Materials folder
30 minutes
Check for
Understanding/Exit
Ticket
Provide evidence of mastery of the day’s
learning objectives.
Homework
Practice Worksheet
Question 4
10 minutes
An ancient sword
made of pure iron is
found. The sword has
reacted with oxygen
gas over the course
of hundreds of years
to form
iron(III)
oxide.
Write the
equation for this
reaction.
See below -Chemical
Reactions
and
Stoichiometry
Practice Worksheet
Youtube video of reaction between sodium and chlorine http://www.youtube.com/watch?v=Mx5JJWI2aaw
PowerPoint Slide Master
Slide 12
Slide 15
Learning Objective:
Write the chemical equation for this reaction:
Blue copper(II) sulfate solution reacts with iron to
form iron(II) sulfate and copper.
 Write chemical equations from word equations.
Reactants:
Copper(II)sulfate and iron
Products:
Iron (II)sulfate and copper
Chemical Equation:
CuSO4
+
Fe →
Slide 13
Chemical Equations
Steps in writing chemical equations:
1. Identify the reactants and the products.
2. Write the formulae (or symbols) of the reactants before the
arrow.
3.Write the formulae (or symbols) of the products after the
arrow.
Slide 14
Chemical Equations
Example:
When magnesium (Mg) is heated, it reacts with oxygen (O2) in
air and burns to produce magnesium oxide (MgO).
Reactants: magnesium (Mg) and oxygen (O2)
Products: magnesium oxide (MgO)
Chemical Equation: Mg + O2  MgO
Slide 16
FeSO4
+
Cu
Write the chemical equation for each reaction:
1. Silver oxide decomposes into silver and oxygen
gas when heated.
2. Ethanol (C2H5OH) burns completely by reacting
with oxygen in air. Carbon dioxide and water
vapor are produced.
3. Aluminum bromide is produced when aluminum
reacts with bromine.
Day 11 (Essential Question(s): ____________________________________________________________________________________)
Learning Objective:

Classify 5 Different Types of Chemical Reactions
Standard 3a
Activity
Demo of single
displacement
double displacement
reactions
Objective
Classify 5 different
types of chemical
reactions starting with
single displacement
Materials Required
(17)
2MCuCl2, Al
1MBaCl2, 1M Na2SO4
Time
5 minutes
Demo double
displacement reactions
Differentiate between
single and double
displacements
1M BaCl2,
1M Na2SO4
5 minutes
Introduce
synthesis,decomposition
& combustion
Check for understanding
Notes
Observe reaction, understand
mechanism of different types of
reactions :
1. Observe reaction,
2. write balanced equation,
3. Predict products and write
balanced equations
Practice classify , predict products,
and write balanced equations
(18) (18A) (18B)
(18C)
Students write balanced
equations on board and Class
discussion
HW: Prelab for “Chemical
Activities of metals
PowerPoint Slide Master
Slide 17
Slide 20
Objective:
Predicting Products
 Predict the products of common chemical
C6H14
+ O2 →
_________
reactions.
 What type of reaction will most likely occur
between the 2 reactants?
 What are the products?
C6H14
Slide 18
+ O2 →
CO2 + H2O
Slide 21
Predicting Products
Al + HCl →
Predict the products of the reaction:
_________
 What type of reaction will most likely occur
between the 2 reactants?
 What are the products?
Al + HCl → AlCl3 + H2
Slide 19
Predicting Products
CuCl2(aq) + Pb(NO3)2 (aq) →
_________
 What type of reaction will most likely occur between the 2
reactants?
 What are the products?
CuCl2(aq) + Pb(NO3)2 (aq) → Cu(NO3)2 (aq) + PbCl2 (s)
1.
NaOH(aq) + FeCl3(aq) →
2.
Zn(s) + HNO3(aq) →
3.
C4H10(g) + O2(g) →
4.
N2(g) + H2(g) →
5.
KBr(aq) + Cl2(g) →
Day 12(Essential Question(s): ____________________________________________________________________________________)
(Version 1)
Learning Objectives:


Compare chemical activities of Cu, Zn, Mg and Ag
Predict products of reaction, write balanced chemical equations
Standard 3a
Activity
Chemical Activity of
Metals Lab
Objective
Observe single
displacement reactions
of Cu,Zn,Mg,Ag
Write balanced
equations of all reacted
reactions
Materials Required
(18)(18A)(18B)(18C)
Time
56 minutes
Notes
Develop critical thinking
skill and based on
results to predict
reactions and write
balanced chem.
equations
Day 12 (Essential Question(s): ____________________________________________________________________________________)
(Version 2)
Learning Objective:
Predict the products of common chemical reactions.
Standard 3a
Activity
Warm Up
Objective
Students
predict
whether mass will
increase, decrease or
decrease
in
the
reaction
between
Alkaseltzer and water in
a sealed ziplock bag.
Materials Required
Time
10 minutes
Notes
To give students hints
on how to write the
experimental
procedure, provide
them a list of
materials they may
use.
Per group:
1 ziplock bag
1 Alkaseltzer tablet
cup with water
triple beam balance
10 minutes
Procedure:
(1) Place ¼ cup water
in ziplock bag.
(2) Place ziplock bag
with water and an
Alkaseltzer tablet on
balance pan. Record
mass.
(3) Add Alkaseltzer
tablet to the water
and immediately seal
the bag. Weigh again.
Chemical Reactions and
Stoichiometry
PowerPoint- see
Supplementary
Materials folder
10 minutes
They will also write a
simple procedure for
testing their prediction.
Alkaseltzer Mini-Lab
Direct Instruction: Law
of Conservation of
Mass and Balancing
Equations
Guided Practice on
Balancing Equations
Students test
prediction.
their
20 minutes
Check for
Understanding/Exit
Ticket
1) 24 grams of carbon
completely reacts with
64 grams of oxygen gas.
What mass of carbon
dioxide is produced?
C + O2 
CO2
24g 64g
?
2) Balance the following
equation:
Al + Fe2O3 → Al2O3 + Fe
Homework
Practice
Worksheet,
Questions 6-7
See attached
Chemical Reactions
and Stoichiometry
Practice Worksheet
PowerPoint Slide Master
Slide 22
Slide 26
Objectives:
Balanced Chemical Equations
Balanced Equation – the number of atoms of each element is
equal on both sides of the equation
 Recognize that chemical reactions are governed
by the Law of Conservation of Mass.
How to balance equations:
1. Count the number of atoms of each element.
2. Use coefficients to make the number of atoms of each
element equal.
3. DO NOT change any of the subscripts.
 Balance chemical equations.
Slide 23
Slide 27
Law of Conservation of Mass
Balancing Chemical Equations
Example 1:
Mg
+ 2 HCl
Reactants:
Mg – 1
H–1 X2=2
Cl – 1 X 2 = 2
Burning Magnesium Metal in an Open Container
→
MgCl2 +
Products:
Mg – 1
H–2
Cl - 2
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 77
Slide 24
Law of Conservation of Mass
Slide 28
Balancing Chemical Equations
Example 2:
+ 2 H2O → 2 NaOH+
H2
2 Na
Reactants:
Products:
Na – 1 X 2 = 2
Na – 1 X 2 = 2
H–2 X2=4
H–1X2+2 =4
O–1 X2=2
O-1 X2=2
Burning Magnesium Metal in a Closed Container
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 77
Slide 25
Law of Conservation of Mass
The total mass of reactants is equal to the total mass of the
products.
Matter is neither created nor destroyed in a chemical
reaction.
2Mg
48 g
Slide 30
+
O2 →
32 g
2MgO
?
Closure: Write-Pair-Share
1.In your own words, describe how a chemical equation is
balanced.
2. Share your answer with your group mates.
3. Make sure that everyone in the group has the correct answer
to the question.
4. If your group is chosen and is able to give the correct answer,
you earn 3 extra credit points.
Slide 29
Balancing Chemical Equations
Balance the following equations:
1. Na
+
Cl2
→
NaCl
2. Fe
+
O2
→
Fe2O3
3. Zn
+
HCl
→
ZnCl2 + H2
4. KNO3
→
KNO2
+
O2
H2
Day 13 (Essential Question(s): ____________________________________________________________________________________)
Learning Objectives:
 Balance chemical equations.
Standard 3a
Activity
Warm Up
Objective
Review balancing of
equations.
Materials Required
Time
5 minutes
Independent Practice This will allow
– Balancing Equations students to practice
further and master
the skill of balancing
equations
See Pogil-Balancing
Equations (pdf file)
35 minutes
Youtube video of the
reaction – see link
below.
10 minutes
Check for
Understanding/Exit
Ticket
Students write a
balanced chemical
equation for the
reaction between
iron and sulfur.
Notes
KClO3 → KCl + O2
Is the above equation
balanced? Why or why
not? Balance the
equation if it’s not.
It is suggested that this
activity be done in small
groups so students will
have a chance to discuss
and clarify concepts and
to check each other’s
work.
Demo – see Activity
Sheet below
You may show a
video of the reaction
or demonstrate it
yourself.
Reaction of Iron and Sulfur Video:
http://www.youtube.com/watch?v=A5H6DVe5FAI&feature=related&safety_mode=true&persist_safety
_mode=1
Day 14 & 15
Thanksgiving Holiday
Day 16 (Essential Question(s): ____________________________________________________________________________________)
Learning Objectives:



Identify the type of chemical reaction.
Predict the products of common chemical reactions.
Write balanced chemical equations.
Standard 3a
Activity
Warm Up
Objective
Go over the objective,
procedure and safety
precautions of the lab.
Materials Required
Time
10 minutes
Rotational Lab Stations
– Types of Reactions
In this lab activity,
students will put
together and apply the
skills they have been
practicing the past
several days –
classifying reactions,
predicting products and
writing balanced
equations.
Check students’
completed activity
sheet.
Practice Worksheet
Question 5
See the materials listed
on the Activity Sheet –
see below
40 minutes
Check for
Understanding/Exit
Ticket
Homework
5 minutes
See above - Chemical
Reactions
and
Stoichiometry Practice
Worksheet
Notes
Lab: Types of Reactions
Purpose: Predict products of a reaction and write chemical equations.
Lab Station 1:
Reaction between calcium carbonate and dilute hydrochloric acid
Prediction: What type of reaction will most likely occur? What products will most likely form?
Explain.
_____________________________________________________________________________
_____________________________________________________________________________
Materials: 50-mL beaker, tiny scoop or spatula, dropper, dilute hydrochloric acid solution,
solid calcium carbonate, waste container
Procedure: Place a small sample of calcium carbonate in the beaker. Add drops of
hydrochloric acid. Observe. Write an equation for the reaction. Dispose of the used chemicals in
the waste container and clean the beaker.
Observation: ______________________________________________________________
Chemical Equation: _________________________________________________________
Lab Station 2:
Reaction between solutions of sodium hydroxide and iron(III) nitrate
Prediction: What type of reaction will most likely occur? What products will most likely form?
Explain.
_____________________________________________________________________________
_____________________________________________________________________________
Materials: test tube, 2 droppers, sodium hydroxide solution, iron(III) nitrate solution, waste
container
Procedure: Mix ten drops of sodium hydroxide solution with ten drops of iron(III) nitrate
solution. Observe. Write an equation for the reaction. Dispose of the used chemicals in the
waste container and clean the beaker.
Observation: ______________________________________________________________
Chemical Equation: _________________________________________________________
Lab Station 3:
Reaction between iron and copper(II) sulfate solution
Prediction: What type of reaction will most likely occur? What products will most likely form?
Explain.
_____________________________________________________________________________
_____________________________________________________________________________
Materials: small test tube, small iron nail, copper(II) sulfate solution, sandpaper, waste
container
Procedure: Half-fill a small test tube with copper(II) sulfate. Place the iron nail in the solution.
Observe. Write an equation for the reaction. Dispose of the used solution in the waste container.
Use sand paper to remove the copper that adheres to the surface of iron nail.
Observation: ______________________________________________________________
Chemical Equation: _________________________________________________________
Lab Station 4:
Burning ethanol (C2H5OH)
Prediction: What type of reaction is this? What products will most likely form?
_____________________________________________________________________________
_____________________________________________________________________________
Materials: dollar bill, large beaker with 50% ethanol solution, large beaker with water, tongs,
matches, paper towel
Procedure: Holding the dollar bill with a pair of tongs, dip it in the beaker of ethanol solution.
With a match, light the dollar bill. Burn the ethanol but not the dollar bill. To prevent dollar bill
from burning, dip it in the beaker of water. (If you burn the dollar bill, you have to pay for it!).
Dry the dollar bill for the next group to use. Write down observations and chemical equation.
Observation: ______________________________________________________________
Chemical Equation: _________________________________________________________
Lab Station 5:
Removing water from copper(II) sulfate pentahydrate, CuSO4∙5H2O
Prediction: What type of reaction is this? What products will most likely form?
_____________________________________________________________________________
_____________________________________________________________________________
Materials: test tube, small scoop or spatula, Bunsen burner, dropper, copper(II) sulfate
pentahydrate, beaker of water
Procedure: Place a tiny sample of solid copper(II) sulfate pentahydrate in a test tube. Take note
of the color. Gently heat the test tube until the solid changes color. Cool down the test tube and
add a few drops of water. Write down observations and chemical equation.
Observation: ______________________________________________________________
Chemical Equation: _________________________________________________________
Day 17 (Essential Question(s): ____________________________________________________________________________________)
Learning Objective:
 Convert moles of reactants to moles of products and vice versa.
Standard 3a, e
Activity
Warm Up
Objective
Students interpret a
recipe and relate it to
interpreting balanced
chemical equations.
Materials Required
Time
10 minutes
Notes
1) 1 bun + 2 patties + 2 cheese
→ 1 cheeseburger
How many patties are needed
to make 5 cheeseburgers?
2) 2H2 + O2 →2H2O
Direct Instruction:
Mole-Mole Conversion
Show students how the
conversion is done
through dimensional
analysis.
Chemical Reactions and
Stoichiometry
PowerPoint- see
Supplementary
Materials folder
Guided Practice
Check for
Understanding/Exit
Ticket
15 minutes
20 minutes
Provide evidence of
mastery of the day’s
learning objective.
10 minutes
How many moles of hydrogen
are needed to make 4 moles of
water?
Emphasize that mole ratio
(conversion factor) is based on
balanced equation. The
coefficients indicate number of
moles.
Write steps on board so
students can refer to them as
they do the practice:
1. Identify known and
unknown.
2. Write possible conversion
factors.
3. Set up equation using
appropriate factors (do the
known units cancel)?
4. Check answer. (sig figs and
unit?)
True or False:
In the reaction shown below, it
takes 1.25 moles of N2 to
produce 2.50 moles of NH3.
N2 + 3H2 → 2NH3
Support your answer with a
calculation.
Homework
Practice Worksheet,
Question 8, 1-d
See above - Chemical
Reactions and
Stoichiometry Practice
Worksheet
PowerPoint Slide Master
Slide 31
Slide 35
Learning Objective:
Mole-Mole Conversion
2H2 +
 Convert moles of reactants to moles of products
O2
→
2H2O
Ex.1 : How many moles of water can be produced from 3.50 moles of
hydrogen?
and vice versa.
Given: 3.50 mol H2
Unknown: mol H2O
Possible conversion factors:
2mol H2
2 mol H2O
Derived from
balanced equation
2 mol H2O
2mol H2
Equation:
moles of H2O
Slide 32
Slide 36
Stoichiometry
x
+
O2
→
Given: 5.0 mol O2
balanced equations.
Equation:
moles of H2
Slide 37
3. Set up equation using appropriate conversion factor(s).
x
mole of unknown
mole of known
Do the known units cancel?
4. Check answer. Sig figs? Units?
Mole-Mole Conversion
2H2 +
2 moles
O2
→
1 mole
x
2 mol H2
1 mol of O2
=
10. moles H2
+
2 HCl
→
MgCl2
+
H2
1. How many moles of magnesium are needed to produce
0.500 moles of magnesium chloride?
2.Write possible conversion factors.
Slide 34
= 5.0 mol O2
Practice Problems:
Mg
mole unknown = mole of known
2mol H2
1 mol O2
1 mol O2
2mol H2
- Stoichiometric calculations are based on
1.Identify given and unknown.
2H2O
Unknown: mol H2
Possible conversion factors:
moles, mass or volume (gases).
Solving Stoichiometric Problems
= 3.50 mol H2O
Ex. 2: How many moles of hydrogen are needed to react with 5.0 moles of
oxygen?
- Amount is usually expressed in number of
Slide 33
2 mol H2O
2 mol of H2
Mole-Mole Conversion
2H2
-Stoichiometry is the calculation of the amount of
reactants and products in a chemical reaction.
= 3.50 mol H2
2H2O
2 moles
Coefficient – indicates number of moles
1.How many moles of oxygen are needed to produce 2 moles of water?
Answer: 1 mole of oxygen
2. How many moles of water can be produced from 4 moles of hydrogen?
Answer: 4 moles of water
3. How many moles of hydrogen is needed to react with 2 moles of oxygen?
Answer: 4 moles of hydrogen
2. How many moles of hydrogen gas can be produced from 6
moles of magnesium?
Day 18 (Essential Question(s): ____________________________________________________________________________________)
Learning Objective:
 Do mass-mole conversions of reactants and products.
Standard 3a, e
Activity
Warm Up
Objective
This is a review of molar
mass to prepare students
to do mass-mole
conversions.
Materials Required
Time
10 minutes
Notes
What is the mass of a mole
of:
1) N2H4
2) Ca(NO3)2
Direction Instruction –
Mole-Mass Conversion
Guided Practice
Check for
Understanding/Exit
Ticket
Show students how the Chemical Reactions and
conversion is done
Stoichiometry
through dimensional
PowerPoint- see
analysis.
Supplementary
Materials folder
15 minutes
Emphasize: Molar mass is
used to do mole-mass
conversion.
20 minutes
Write steps on board so
students can refer to
them as they do the
practice:
1. Identify known and
unknown.
2. Write possible
conversion factors.
3. Set up equation using
appropriate factors (do
the known units
cancel)?
4. Check answer. (sig
figs and unit?)
Provide evidence of
mastery of the day’s
learning objective.
10 minutes
2KClO3 → 2KCl + 3O2
How many grams of KCl
can be produced from 0.50
moles of KCl?
(a) Name the given and the
unknown in the problem.
(b) What possible
conversion factors can you
use to solve the problem?
(b) Show how the equation
should be set up.
Homework
Practice Worksheet,
Question 9, 1-d
See above - Chemical
Reactions and
Stoichiometry Practice
Worksheet
PowerPoint Slide Master
Slide 38
Mole-Mass or Mass-Mole Conversion
Given
Molar mass
mol B = mass A x
mass B = mol A x
Slide 39
Mole ratio from
balanced equation
1 mol A
x mol B
mass A
mol A
mol B
mol A
x mass B
1 mol B
Mole-Mass or Mass-Mole Conversion
2H2
+
O2
→
2H2O
Ex.1 : What is the mass of oxygen that is needed to produce 4.0 moles of water?
Given: 4.0 mol H2O
Unknown: g of O2
Possible conversion factors: 1 mol O2
2 mol H2O
Equation:
mass of O2 = 4.0 mol H2O x
Slide 40
32g O2
1 mol O2
1 mol O2 x 32g O2 = 64 g O2
2 mol H2O
1 mol O2
Practice Problems:
Mg +
2 HCl
→
MgCl2
+
H2
1.How many moles of magnesium are needed to
form 47 grams of magnesium chloride?
2. How many grams of magnesium are needed to
produce 4.5 moles of hydrogen?
Day 19 (Essential Question(s): ____________________________________________________________________________________)
Learning Objective:
 Do mass-mass conversions of reactants and products.
Standard 3a, e
Activity
Warm Up
Objective
Review mole-mass
conversion.
Materials Required
Time
10 minutes
Direct Instruction –
Mass-Mass Conversion
Show students how the
conversion is done
through dimensional
analysis.
Chemical Reactions
and Stoichiometry
PowerPoint- see
Supplementary
Materials folder
15 minutes
Guided Practice
Check for
Understanding/Exit
Ticket
Provide evidence of
mastery of the day’s
learning objectives.
Notes
N2 + 3H2 → 2NH3
How many grams of
ammonia can be produced
from 1.50 moles of
hydrogen gas?
20 minutes
Write steps on board so
students can refer to
them as they do the
practice:
1. Identify known and
unknown.
2. Write possible
conversion factors.
3. Set up equation using
appropriate factors (do
the known units cancel)?
4. Check answer. (sig figs
and unit?)
10 minutes
Mg + CuSO4 → MgSO4 + Cu
How many grams of Cu can
be produced when 5.00 g of
Mg reacts completely with
CuSO4?
(a) What are the steps in
solving the above problem?
Give the correct conversion
factor for each step.
(b) Set up an equation that
shows the conversion
factors you listed in (a).
Homework
Practice Worksheet,
Question 10, 1-d
See above - Chemical
Reactions and
Stoichiometry Practice
Worksheet
PowerPoint Slide Master
Slide 41
Mass- Mass Conversion
2H2
2(2g) = 4g
+
O2
1(32g) = 32 g
→
2H2O
2 (18g) = 36g
1. How many grams of hydrogen are needed to produce 36 g of
water?
Answer: 4g hydrogen
2. How many grams of water can be produced from 32 g of
oxygen?
Answer: 36 g water
3. What is the mass of oxygen that is needed to react with 8 g of
hydrogen?
Answer: 2(32g) = 64 g
Slide 42
Mass -Mass Conversion
mass B = mass of A x 1 mole A
mass A
2H2
+
O2
x mole B
mole A
→
x mass B
1 mole B
2H2O
Ex.1 : What is the mass of oxygen that is needed to produce 18 g of
water?
Given: 18 g of H2O
Unknown: g of O2
Possible conversion factors: 1 mol H2O
18 g H2O
Equation:
mass of O2 = 18 g H2O x 1 mol H2O
18 g H2O
Slide 43
1 mol O2
32g O2
2mol H2O 1 mol O2
x 1 mol O2
x 32 g O2 = 16 g O2
2 mol H2O
1 mol O2
Practice Problems:
Mg
+
2 HCl
→
MgCl2
+
H2
1. How many grams of magnesium are needed to produce 6g of
hydrogen?
2. How many grams of magnesium chloride can be produced
from 54 g magnesium?
Day 20 (Essential Question(s): ____________________________________________________________________________________)
Learning Objective:
 Convert mass to mole to molar volume of gaseous reactants and products at STP.
Standard 3a, d, e
Activity
Warm Up
Objective
Review mass-mole
conversion.
Materials Required
Time
5 minutes
Notes
Zn + HCl → ZnCl2 + H2
How many moles of
hydrogen gas can be
prepared from the
reaction of 2g of Zn
with excess HCl?
Demo – Hydrogen
Balloon Explosion
Hydrogen gas will be
generated from the reaction
between Zn (or Mg) and HCl.
Given the mass of Zn used in
the reaction, guide the
students in figuring out the
volume of the hydrogen gas
produced, assuming
standard conditions for
pressure and temperature.
From the balanced equation
for the reaction between H2
and O2, let the students
figure out the volume of O2
needed to completely burn
the H2 gas.
Guided Practice
Check for
Understanding/Exit
Ticket
See Activity Sheet below
25 minutes
Chemical Reactions and
Stoichiometry
PowerPoint- see
Supplementary
Materials folder
20 minutes
2C4H10 + 13O2 → 8CO2 +
10H2O
5 minutes
If 0.33 moles of butane
(C4H10) are burned, how
many liters of carbon dioxide
would be produced at STP?
Homework
Practice Worksheet,
Question 11 a-d
See above - Chemical
Reactions
and
Stoichiometry
Practice
Worksheet
Be aware of the safety
precautions that must
be observed in
demonstrating the
ignition of H2 gas. See
Activity Sheet below.
PowerPoint Slide Master
Slide 44
Slide 47
Learning Objectives:
Practice Problems:
Mg
 Perform mass-mole-volume conversion at STP.
2 HCl
→
MgCl2
+
H2
1. How many grams of magnesium are needed to produce 11.2
L of hydrogen gas at STP?
 Perform volume-volume conversion of gaseous
2. How many liters of hydrogen gas at STP may be produced
from the reaction of 15.0 g of magnesium with excess
hydrochloric acid?
reactants and products at STP.
Slide 45
+
Slide 48
Molar Volume
Volume -Volume Conversion (at STP)
L of B = L of A x
 Avogadro’s Principle: Equal volumes of gases at
the same temperature and pressure contain the
same number of particles.
2H2
+
O2 →
mol B
mol A
2H2O
Ex.1 : How many liters of oxygen gas are needed to completely react with 13.5 L of
hydrogen gas at STP?
 At STP (Standard Temperature and Pressure),
Given: 13.5 L H2
1 mole of any gas occupies a volume of 22.4 L.
Unknown: L of O2
Possible conversion factors:
1mol O2
2 mol H2
Equation:
L of O2 = 13.5
L H2 x 1mol O2
=
6.75 L O2
2 mol H2
Slide 46
Mass-Mole-Volume Conversion (at STP)
L of B = mass of A x
2 KClO3
1 mol A
mass A
→
x mol B
mol A
x
22.4 L B
1 mol B
2KCl + 3O2
Ex.1 : How many liters of oxygen gas are produced when 30.0 g of potassium chlorate
decomposes at STP?
Given: 30.0 g KClO3
Unknown: L of O2
Possible conversion factors: 1 mol KClO3
3mol O2
22.4 L O2
122.5 g KClO3
2mol KClO3 1 mol O2
Equation:
L of O2 = 30.0 g KClO3 x 1 mol KClO3 x 3 mol O2
x 22.4 L O2 = 8.23 L O2
122.5 g KClO3 2 mol KClO3
mol O2
Slide 49
Practice Problems:
N2(g) +
3H2(g)
→
2 NH3(g)
1. How many liters of hydrogen gas are needed to completely
react with 40.0 L of nitrogen gas at STP?
2. How many liters of ammonia gas may be produced when
50.0 L of hydrogen gas react with excess nitrogen gas at STP?
Day 22 (Essential Question(s): ____________________________________________________________________________________)
Learning Objective:
 Distinguish between limiting and excess reactants.
Standard 3f(not tested on CST)
Activity
Warm Up
Objective
1 bun + 2 patties + 2
cheese slices → 1
cheeseburger
Materials Required
Time
10 minutes
Chemical Reactions
and Stoichiometry
PowerPoint- see
Supplementary
Materials folder
15 minutes
If there are 8 buns, 12
patties and 12 cheese
slices available, how
many cheeseburgers
can we make?
Direction Instruction –
Solving limiting and
excess reactant
problems
Guided Practice
Check for
Understanding/Exit
Ticket
Ingredients needed for
making pancakes: 1 cup
flour, ½ cup milk, 1 egg
Ingredients available on
hand: 2 cup flour, 2
cups milk, 2 eggs
10 minutes
1.Which of the
available ingredient(s)
is/are in excess?
2. Which of the
available ingredient(s)
limit(s) the amount of
pancakes that can be
made?
2H2 + O2 → 2H2O
3. If 1.50 moles H2 and
0.50 moles O2 react,
will both reactants be
completely consumed?
If not, name the excess
reactant.
Homework
Practice Worksheet –
Question 12 a-b
20 minutes
See above - Chemical
Reactions and
Stoichiometry Practice
Worksheet
Notes
Using food recipes as
examples is an
engaging way to
introduce the difficult
concepts of limiting and
excess reactants. You
may have students
suggest their own
recipes and have them
come up with similar
questions as on the
warm up.
PowerPoint Slide Master
Slide 50
Slide 53
Learning Objective:
Limiting and Excess Reactants
2H2
 Distinguish between limiting and excess
+
O2
→
2H2O
Ex.1: 6.0 g of H2 and 60.g of O2 are made to react.
reactants.
(a)
Is there a reactant present in excess? If there is, how many grams
of this reactant is left unreacted?
(b)
How many grams of water are produced from the reaction?
mass of O2 = 6.0 g H2 x 1 mol H2 x 1 mol O2 x 32g O2 = 48 g O2
2 g H2
2 mol H2
1 mol O2
Only 48 g of O2 is needed to completely react, so it is an excess reactant. 12 g of
O2 is left over.
Slide 51
Limiting and Excess Reactants
Slide 54
Limiting and Excess Reactants
2H2
1 bun + 2 patties + 2 cheese slices → double
cheeseburger
+
O2
→
2H2O
Ex.1: 6.0 g of H2 and 60.g of O2 are made to react.
If there are 5 buns, 8 patties and 6 cheese slices
available, how many double cheeseburgers can be
made?
(a)
Is there a reactant present in excess? If there is, how many grams
is left unreacted?
(b)
How many grams of water are produced from the reaction?
H2 is the limiting reactant; it determines the amount of water produced.
Which ingredient is completely used up?
mass of H2O = 6.0 g H2 x 1 mol H2 x 2 mol H2O x 18 g H2O = 54g H2O
2 g H2
2 mol H2 1 mol H2O
Which ingredient is left over?
Slide 52
Limiting and Excess Reactants
Slide 55
Practice Problems:
2Al
Limiting Reactant – completely used up; limits the
amount of product
Excess Reactant – not completely used up, “left over”
+
3Br2
→
2AlBr3
20 g aluminum and 100.0 g bromine were made to react.
1. What is the limiting reactant in the reaction?
2. How much of the excess reactant is left over after the
reaction?
3. How many grams of aluminum bromide is produced from
the reaction?
Day 23 (Essential Question(s): ____________________________________________________________________________________)
Learning Objective:
 Calculate the theoretical yield and percent yield of a reaction.
Standard 3f(not tested on CST)
Activity
Warm Up
Description/Details
Review mass-mass
conversion.
Materials Required
Time
10 minutes
Students offer possible
explanations for why
actual yield is usually
less than theoretical
yield.
Direct Instruction and
Guided Practice –
Solving Theoretical and
Percent Yield Problems
Chemists often get less than the
maximum amount of product that
they expect from a reaction. Why
do you think this is so?
Chemical Reactions
and Stoichiometry
PowerPoint- see
Supplementary
Materials folder
Pre-Lab Discussion –
Preparation of Salt
Go over objective,
procedure and safety
precautions of the lab.
Check for
Understanding/Exit
Ticket
Provide evidence of
mastery of the day’s
learning objective.
Notes
2Na + Cl2 → 2NaCl
What is the maximum amount (in
grams) of NaCl that can be
produced if 56.0 g sodium reacts
with excess chlorine?
30 minutes
10 minutes
The lab may be modified by
giving only general directions
and having the students figure
out the procedure on their own.
For example, you may instruct
them to prepare sodium
chloride from 1 g of baking soda
and excess HCl. Have them
figure out ways to minimize
errors and maximize the yield of
the reaction. Their procedure
should also include safety
precautions. If this modification
is made, 1 more period is
needed for students to work
with their group mates in
formulating the procedure.
NaHCO3 + HCl → NaCl + H2O + CO2
How many grams of NaCl may be
produced if 8.40g NaHCO3 reacts
completely with excess HCl
solution?
What is the percent yield of the
reaction, if only 4.0 g NaCl is
actually produced?
Homework
Answer Pre-lab
Questions –
Preparation of Salt
PowerPoint Slide Master
Slide 56
Slide 59
Learning Objective:
Sample Problem 2
Mg
+
2 HCl
→
MgCl2
+
H2
 Determine the theoretical and percent yield of a
reaction.
1.How many grams of hydrogen are formed from 24 g
magnesium?
Answer: 2 g
2. If only 1g of hydrogen is actually produced from 24 g
magnesium, what is the % yield of the reaction?
Answer:
Slide 57
Theoretical and Percent Yield
Theoretical Yield – amount of product formed when all of the
reactants are completely used up
Actual Yield – amount of product actually formed in a
reaction
Usually: Actual Yield < Theoretical Yield
Percent Yield – indicates how well a reaction comes to
completion
Percent Yield =
Slide 58
Actual
Theoretical
x 100
Sample Problem 1
2H2
2(2g) = 4g
+
O2
1(32g) = 32 g
→
2H2O
2 (18g) = 36g
1. How many grams of water can be produced from 32 g of
oxygen?
Answer: 36 g water (theoretical yield)
2. If only 27 g of water is actually produced from 32 g of
oxygen, what is the % yield of the reaction?
Answer: % yield = 27 x 100 = 75%
36
Slide 60
% yield = 1g x 100 = 50%
2g
Check for Understanding
2KClO3
→
2KCl
+
3O2
What is the % yield of the above reaction if only 45 g of oxygen
is produced from 122 g of potassium chlorate?
Day 24 (Essential Question(s): ____________________________________________________________________________________)
Learning Objective:
 Calculate the theoretical yield and percent yield of a reaction.
Standard 3f(optional)
Activity
Warm Up
Lab – Preparation of
Salt
Check for
Understanding/Exit
Ticket
Homework
Description/Details
Students review the
procedure and safety
precautions with
their group mates.
In this lab activity,
students will apply
the skills they have
been practicing the
previous days – doing
mass – mass
conversion and
calculating
theoretical and
percent yield.
Check data and
calculations.
Answer analysis
questions and write
conclusion.
Materials Required
See Activity Sheet
below for list of
materials
Time
5 minutes
45 minutes
5 minutes
Notes
Lab: Preparation of Salt
Overview and Purpose:
The percent yield of a reaction will be determined in this experiment by comparing the actual
mass of product formed from the reaction with the theoretical (or expected) mass.
x 100
The reaction to be studied is the double displacement reaction between sodium bicarbonate
(NaHCO3), commonly known as baking soda, and hydrochloric acid (HCl). The reaction produces table
salt or sodium chloride (NaCl) and carbonic acid (H2CO3). The carbonic acid readily decomposes to water
and carbon dioxide as shown by the equations below.
NaHCO3 + HCl → NaCl + H2CO3
NaHCO3 + HCl → NaCl + H2O + CO2
A known mass of NaHCO3 will be reacted completely with an excess of HCl. From the actual
mass of sodium chloride produced and the calculated theoretical yield, the percent yield of the reaction
can be determined.
Pre-Lab Questions: (Write your answers on a separate sheet of paper. Turn in Pre-Lab paper the day
before the lab)
A sample of silver oxide (Ag2O) was heated several times to drive out all of the oxygen, leaving
behind silver, according to the reaction,
2Ag2O → 4Ag + O2
The following data were collected.
Mass of crucible (g)
20.5552
Mass of crucible + Ag2O (g)
22.5535
Mass of crucible + remaining solid after first heating (g)
22.38220
Mass of beaker + remaining solid after second heating (g)
22.1621
Mass of beaker + remaining solid after third heating (g)
22.1621
(1) How many grams of silver oxide (Ag2O) was used in the reaction?
(2) How many grams of the solid (Ag) remained after the third heating?
(3) How many grams of Ag are expected to be produced from the grams of silver oxide used in the
reaction?
(4) Calculate the percent yield of the reaction.
Materials: 250-mL beaker, weighing paper, spoon, dropper, hot plate, baking soda (NaHCO3) and
hydrochloric acid (HCl).
Procedure:
1. Weigh an empty beaker.
2. Place about 1 g of NaHCO3 into the beaker. Record the mass of the beaker and the NaHCO3.
3. Add HCl to NaHCO3 drop by drop. When the reaction mixture stops fizzing all of the NaHCO3
has been reacted.
4. Gently heat the beaker on a hot plate until all the liquid has evaporated.
5. Carefully take the beaker off the hot plate. Allow it to cool and take the mass of the beaker and
the salt that is left in the beaker.
6. Repeat steps 4-5 until no more decrease in mass is observed.
Data:
Mass of empty beaker (g)
Mass of beaker + NaHCO3 (g)
Mass of beaker + NaCl after first heating (g)
Mass of beaker + NaCl after second heating (g)
Mass of beaker + NaCl after third heating (g)
Calculation:
1. Calculate the mass of the NaHCO3 used in the reaction.
2. Calculate the mass of NaCl actually produced from the reaction.
3. Calculate the theoretical yield of the reaction, that is, the mass in grams of NaCl that is expected to be
produced from mass of NaHCO3 used in the reaction.
4. Calculate the percent yield of the reaction
Analysis:
Why was the NaCl produced from the reaction heated several times? How does the actual yield compare
to the theoretical yield? List 3 possible sources of error in the experiment and explain how these errors
affected the result. (Be specific!)
Conclusion:
What was the purpose of the experiment? Was this purpose achieved? What have you learned from the
experiment?
Day 25, 26, 27 Review – Benchmark Assessment and Finals
PowerPoints are attached. Review may be done using a Cooperative Learning strategy called
Numbered-Heads-Together. Students working in groups of 3 figure out answer to each
question together. Group member whose number is called writes the answer on white board
for teacher to check.
Slide 1
Slide 7
NH3 + HCl → NH4Cl
How many grams of NH3 are
needed to produce 25 g of
ammonium chloride?
Test Review
Chemical Reactions and
Stoichiometry
Slide 2
8.0 g NH3
Slide 8
2NH3 → N2 + 3H2
This test includes:
1. identifying types of reactions
2. writing and balancing of equations
3. calculation of moles, mass and volume of
reactants and products.
4. calculation of the % yield of a reaction.
How many moles of NH3 are
needed to produce 2.0 g of
hydrogen?
0.67 moles NH3
Slide 3
Slide 9
CaCO3 → CaO + CO2
Propane reacts with oxygen to produce water and
carbon dioxide based on the equation:
Which is true of the above reaction?
C3H8 +
A. Calcium oxide and carbon dioxide are reactants.
B. The reaction involves the decomposition of calcium
carbonate.
C. The reaction involves the combination of calcium
oxide and carbon dioxide.
D. The reaction involves the combustion of calcium
carbonate.
5O2
→
4H2O +
3CO2
What volume of propane at STP is needed to
produce 6.0 moles of CO2?
44.8 L
B
Slide 4
Slide
10
Which is a single displacement reaction?
A. NH4NO2 → N2 + 2H2O
B. Na2SO4 + Ba(NO3)2 → BaSO4 + 2NaNO3
C. Fe + 2HCl → FeCl2 + H2
D. C2H5OH + 3O2 → 2CO2 + 3H2O
Propane reacts with oxygen to produce water and
carbon dioxide based on the equation:
C3H8 +
5O2
→
4H2O +
3CO2
How many grams of oxygen gas are needed to
completely burn 22.4 L of propane at STP?
160 g
C
Slide 5
Slide
2H2O → 2H2 + O2
11
Which of these equations is balanced?
A. NH4NO2 → N2 + H2O
B. NH3 + HCl → NH4Cl
C. 2NH3 → N2 + H2
D. NH4NO3 → N2O + H2O
B
1. How many grams of oxygen can be formed
from the decomposition of 36 g of water?
2. If only 30 g of oxygen is actually produced
from 36 g of water, what is the % yield of the
reaction?
32 g O2, 94%
Slide 6
NH3 + HCl → NH4Cl
How many moles of hydrogen
chloride are needed to produce
0.35 moles of ammonium
chloride?
0.35 moles HCl
Slide
12
4Al+ 3O2 → 2Al2O3
100.0 g of Al and 100.0 g of O2 react. What
are the limiting and excess reactants? How
many grams of Al2O3 are produced?
Limiting: Al
Excess: O2
188.9 g Al2O3
Slide 1
Slide 6
Calculate the % error of a measurement
(0.90 g/mL) if the true value is 1.0 g/mL.
Answer: 10%
Final Exam Review
Chemistry – First Semester
Slide 2
“Matter is made up of atoms. Atoms have tiny positive
centers containing protons and neutrons.”
Slide 7
The above statement is a/an
A. inference
B. hypothesis
C. theory
D. observation
Answer: C
Which of these statements is NOT true about matter and
energy?
A. All matter possesses energy.
B. Matter can be changed into energy and energy can be
changed into matter.
C. Energy can only be transferred from one sample of
matter to another when they are in direct contact with
one another.
D. The total amount of matter and energy in the universe
remains the same; they just change from one form to
another.
Answer: C
Slide 3
In order to become a theory, a hypothesis should be
A. obviously accepted by most people.
B. a fully functional experiment.
C. in alignment with past theories.
D. repeatedly confirmed by experimentation
Slide 8
Answer: D
Which of these is an exothermic process?
A. Melting of ice
B. Combustion (burning) of gasoline
C. Photosynthesis
D. Evaporation of water
Answer: B
Slide 4
Slide 5
What is the density of a substance which
has a volume of 100.0 mL and a mass of
85.5 g?
Answer: 0.855 g/mL
How many significant digits does this
measurement have?
95.50 mL
Answer: 4
Slide 9
Slide 10
An endothermic reaction
A. releases energy to the surroundings.
B. causes a temperature increase in its
surroundings.
C. absorbs energy from its surroundings.
D. produces substances that have a lesser
energy than the starting materials.
Answer: C
Which is true about metals?
A. They are found on the right side of the periodic table.
B. They are poor conductors .
C. They have higher densities compared to non-metals.
D. Most of them are liquids and gases at room
temperature.
Answer: C
Slide 11
A.
B.
C.
D.
The nucleus of an atom
is negatively-charged
accounts for most of the atom’s mass
occupies most of the atom’s volume
contains electrons and protons
Slide 16
When a metal is heated in a flame, the flame turns a
distinctive color. This information can be applied in
the study of stars because
A. star color tells us how far it is from the earth.
B. the color of the star tells us how big it is.
C. the color of the star tell us how old it is.
D. the color spectra of a star show which elements are
present in it.
Answer: D
Answer: B
Slide 12
Give the number of protons, electrons and
neutrons of 158O.
Slide 17
How many valence electrons does nitrogen
have?
Answer: electrons- 8, protons- 8, neutron-7
Slide 13
Which of these is NOT true of the atom?
A. It may be positively or negativelycharged.
B. It contains the same number of protons
and electrons.
C. It is mostly empty space.
D. It has a very dense center.
Answer: 5
Slide 18
Draw the Lewis Dot diagram of carbon
Answer:
C
Answer: A
Slide 14
Isotopes of the same element always have the same
A. atomic number
B. mass number
C. atomic mass
D. number of neutrons
Slide 19
Which element do you expect to have the
same valence electrons as carbon?
A.
B.
C.
D.
Answer: A
calcium
iron
chlorine
silicon
Answer: D
Slide 15
Some isotopes easily break down and emit
radiation. Which type of radiation is the
least penetrating?
A. beta
B. alpha
C. X-ray
D. gamma
Answer: B
Slide 20
Write the electron configuration of
potassium.
Answer: 1s2 2s2 2p6 3s2 3p6 4s1
Slide 21
Slide 22
Slide 23
Which of these is a halogen?
A. helium
B. carbon
C. lithium
D. iodine
Answer: D
To which family does the element krypton
belong?
A. Alkali metal
B. Alkaline-earth metal
C. Transition Metal
D. Noble gases
Answer: D
Which of these is true of noble gases?
A. They have 8 valence electrons except for
helium.
B. They belong to Group 2A.
C. They easily form bonds with other
elements.
D. They have very low ionization energies.
Answer: A
Slide 26
Which element is the least electronegative?
A. Nitrogen
B. Phosphorus
C. Oxygen
D. Sulfur
Answer: B
Slide 27
Which type of chemical bond exists
between a metal and a non-metal?
Ionic
Slide 28
Answer: A
Slide 24
Slide 25
In general, how do atomic masses change
throughout the periodic table of elements?
A. They increase from left to right and top to
bottom.
B. They increase from left to right and bottom
to top.
C. They increase from right to left and top to
bottom.
D. They increase from right to left and bottom
to top.
Answer: A
Why is cobalt (Co) placed before nickel (Ni)
on the periodic table even though its
atomic mass is higher than nickel’s?
A. Cobalt has more electrons.
B. Nickel has a higher density.
C. Cobalt was discovered first.
D. Nickel has one more proton.
Answer: D
Slide 29
Metallic substances are usually good conductors
of electricity because
A. of the loosely bonded electrons that move
around the metal structure
B. of the closely packed atoms
C. of the strong bonds that exist between atoms
D. of the weak bonds that exist between atoms
Answer: A
Slide 30
Which of these is NOT an electrical
conductor?
A. aluminum
B. Alcohol (C2H5OH) solution
C. Potassium chloride (KCl) solution
D. Hydrochloric acid (HCl) solution
Answer: B
Slide 31
Slide 32
Which element can form long chains of
atoms by forming single, double and
triple bonds with itself?
A. oxygen
B. nitrogen
C. carbon
D. hydrogen
Answer: C
What is the correct formula of sodium
bromide?
A. NaBr
B. Na2Br
C. NaBr2
D. Na2Br2
Answer: A
Slide 33
Slide 36
Give the name of
FeSO4
Iron (II) sulfate
Slide 37
What is the correct formula of sodium
bromide?
A. NaBr
B. Na2Br
C. NaBr2
D. Na2Br2
Answer: A
Slide 38
Which of these is an acid?
H2SO4
KOH
K2SO4
What is the chemical formula for
sodium sulfate?
Na2SO4
Answer: H2SO4
Slide 34
Slide 39
What is the chemical formula for
copper (II) nitrate?
Which is a polar covalent bond?
A. N-H
B. As-H
C. O-O
D. Cl-Cl
Cu(NO3)2
A
Slide 35
Slide 40
Give the name of
CaCl2
Write the Lewis dot diagram of
H2S.
Calcium Chloride
H
S
H
Slide 41
Slide 46
How many non-bonding (unshared)
pairs of electrons does sulfur have in
H2S?
Predict the shape of the NH3.
2
Trigonal Pyramidal
Slide 42
Slide 47
These pairs of atoms are covalently bonded.
Arrange them according to increasing polarity.
A. N-O
B. Cl-F
C. C-H
D. C-Cl
How many non-bonding pairs
does sulfur have in H2S?
2
Slide 43
C,D,A,B
Slide 48
Which of these is an alkene?
A. C3H8
B. C4H8
C. C5H8
D. C6H14
What is the shape of the H2S
molecule?
Bent
Slide 44
B
Slide 49
Is the H2S molecule polar or nonpolar?
Polar
B
Slide 45
Slide 50
NH3 + HCl → NH4Cl
Which of these is a polar molecule?
A. CO2
B. SF6
C. CH4
D. PCl3
How many grams of NH3 is needed to
produce 25 g of NH4Cl?
8.0 g of NH3
D
Day 28, 29, 30
Benchmark Assessment 3
Final Exams