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 10 23 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
Objective
Describe and interpret one mole of atoms, formula units (f.u.), molecules as related to the mass of elements and compounds
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(H
2
O)
7 different jars
Pass around classroom for observations
Avogadro’s No. the concept of 1 mole
=6.02X10
23 Use dimensional analysis for Mole Conversion
Convert between moles and particles using conversion factors:
1mole/6.02X1023particles
6.02X1023particles/1mole
(1) Lecture Notes and
Worksheet on Mole to
Particles Conversion
Time Notes
15 minutes 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)
35 minutes Students write answers of every question on board show each step
Class Discussion of student work
Check for
Understanding
1 /2mole of H
2
O= ? molecules
1.2X10
24 molecules of
H
2
O=?mole
5 minutes (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
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
Objective Materials Required
12g Carbon
And Previous day samples
Time Notes
10 minutes Review mole, particles and atomic masses
Power points on conversion of molar mass of elements, ionic molecular, compounds to moles
Check for understanding
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
(13) Diagram: Formula Mass of
H
2
O
Worksheet (2) + Answers (3) on
(4)
(7)
Molar Mass Calculation
Lecture figures on mass to mole conversion
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.
40 minutes
5 minutes
Work on
Practice
Questions
Worksheet(2)
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
Review mole to mass & mole to Particles
Conversions
Convert mass to mole to particles
Objective
Convert moles of elements, ionic compounds and molecular compounds to mass and particles
Summarize Mole Conversion and expand mass to mole to particles conversions
Practical Example of Conversion from mass to fu
Check for understanding
Discuss results based on student work
Materials Required Time Notes
(8) Mass to Mole Power Points (B) 10 minutes 5.50g NaCl=? mole
5.50X10
24 =? fu
(5) Lecture Notes on Mole
Conversion
(5A) In-class mole assignment
(9) 10:2 Worksheet
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

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 Mini-
Lab
(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, H
2
O, electronic scales
Time Notes
45 minutes 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, CuSO
4
∙ x H
2
O, 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 Na
2
CO
3
∙ x H
2
O, where x is the number of moles of H
2
O per mole of Na
2
CO
3
. When a 2.123 g Na
2
CO
3
∙ x H
2
O was heated at 130 o C, all of the water of hydration was lost, leaving 0.787 g Na
2
CO
3
.
(1) Calculate the mass of water lost.
(2) Convert the mass of water lost to moles.
(3) Convert 0.787 g Na
2
CO
3
to moles.
(4) What is the ratio of moles of H
2
O to moles of Na
2
CO
3
?
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 (after1 st heating) (g)
Mass of crucible and anhydrous copper sulfate (after 2 nd 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 (CuSO
4
).
(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 CuSO
4
∙
5 H
2
O. 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 CuSO
4
∙ 5 H
2
O?

Mole Mini-Lab
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, CaCO
3
. What is the molar mass of CaCO
3
?
(c) Convert the mass of CaCO
3
to moles.
5. So how many moles of chalk is your name worth? __________________________
For extra credit, calculate the number of ions of Ca 2+ and CO
3
2 that you used to write your name.

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
2.Molar
Volume
Conversion
Check for understanding
Objective
Review conversion from Mass to
Mole to Particles
Define STP
Define Molar Volume as 22.4L at
STP
Conversion of Volume to Mole to
Mass
Students Recreate Mole Map to enhance Mole Conversion
Materials Required
(11) Mole Map
(6) Mole Conversion
Worksheet
(6A) Answers for (6)
Time Notes
20 minutes Use Mole Map to guide conversion
(14) Volume to Mole and
Mole Map for Review
(6) Mole Conversion
Worksheet
30 minutes Do (6) in class and HW: (9)
10:2 Practice
Problems
(9) 10:2 Practice Worksheet
#23 to #30
5 minutes Draw Mole
Map
Check for X or
/ for mole conversion based on Mole
Map

Day 7
(Essential Question(s): ____________________________________________________________________________________)
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 Test Taking
Skill
Materials Required
(15) and (16) Review
Questions
Similar to questions in
STAR/SAT tests
Day 7
Time
56 minutes
Notes
Students write answers on board and show work
Class discussion of all problems

Slide 1
Slide 2
Slide 3
Slide 4
Slide 5
Slide 6
Test Review
Mole Concept
How many particles are there in
0.25 moles of a substance?
1.5 x 10 23 particles
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 10 21 particles.
B
Slide 7
Slide 8
What is the molecular mass of
C
4
H
10
?
58 amu
What is the mass of 6.02 x 10 23 molecules of CO
2
?
44 g
Which is equal to 45 g of H
2
O?
A. 1 mole
B. 1.5 moles
C. 2.0 moles
D. 2.5 moles
D
Slide 9
What is the molar mass of HNO
3
?
63 g
What is the molar mass of CaCl
2
?
110 g
Slide 10
How many moles are there in
60.0 g of carbon?
What is the mass of 0.75 moles of CaCl
2
?
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 (H
2
O
2
)?
B. What is the molar mass of calcium chloride (CaCl
2
)?
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 (H
2
0)?
B. What is the mass of 1.5 moles of sugar (C
12
H
22
O
11
)?
C. How many formula units (particles) of sodium chloride (NaCl) are in 100.0 g of this substance?

Day 8
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.
O
2
B.
CH
3
OH
C.
SO
2
D.
CaCl
2
2) What is the molar mass of CaCl
2
?
A.
110. g
B.
90.0
C.
75.0 g
D.
70.0 g
3) *How many moles are contained in 9.03 x 10 23 molecules of oxygen gas (O
2
)?
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 H
2
O?
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.
0 o C and 2 atm
B.
100 o C and 1 atm
C.
0 o C and 1 atm
D.
100 o C 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(NO
3
)
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 10 24 nitrate (NO
3
) ions.
10) 9.03 x 10 23 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 MgCl
2
· n H
2
O. She placed a small sample of MgCl
2
· n H
2
O 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 22.347 g
Initial mass of sample and container
Mass of sample and container after first heating
25.825 g
23.982 g
Mass of sample and container after second heating
Mass of sample and container after third heating
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 MgCl
2
that remain in the crucible.
(iv) calculate the mole of MgCl
2
that remain in the crucible.
(v) How many moles of water are lost per mole of MgCl
2
? 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
Objectives
Review the difference between physical and chemical change.
Materials Required
Direct Instruction and
Guided Practice -
Evidences of Chemical
Change, Reactants and
Products, Types of
Reaction
Students identify evidences of a chemical change, identify reactants and products and classify reactions.
Chemical Reactions and
Stoichiometry PowerPoint- see Supplementary
Materials folder
Demo: Mg, HCl solution, paper, matches, Pb(NO
3
)
2 and KI solution, beaker, test tubes, droppers
Time Notes
5 minutes Physical or Chemical
Change?
1. melting of ice
2. rotting of food
3. burning of gasoline
4. evaporation of water
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(NO
3
)
2
+ KI
10 minutes Check for
Understanding/Exit
Ticket
Homework
NaHCO
3
+ CH
3
COOH →
CH
3
COONa + H
2
CO
3
NaHCO
3
+ CH
3
COOH →
CH
3
COONa + H
2
O + CO
2
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?
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 2
Chemical Reactions and
Stoichiometry
Objectives:
 Recognize signs of chemical reactions.
 Recognize the reactants and products in a reaction.
 Identify different types of chemical reactions.
Slide 3
Signs of Chemical Reactions
Evolution of a gas
Slide 5
Signs of Chemical Reactions
Formation of a precipitate
Precipitate –
insoluble solid formed from the reaction between 2 aqueous solutions
Slide 6
Chemical Reaction and Equation
Chemical reaction – a change that forms new substances
Reactants – starting substances
Products – new substances formed
Chemical Equation
Shorthand way of describing chemical reactions
Example:
2H
2
(Reactants)
+ O
2
→
2H
2
O
(Product)
Slide 7
AgNO
3
+ NaCl → AgCl + NaNO
3
What are the reactants in the above reaction?
What are the products?
Slide 4
Signs of Chemical Reactions
Change in intensive properties like color, odor, density
Release or absorption of energy
Slide 9
Types of Chemical Reactions
3. Single Displacement
2 reactants → 2 products
(active element and compound)
Example: Mg + HCl → MgCl
2
+ H
2
4. Double Displacement
2 reactants → 2 products
(2 aqueous solutions)
Example: KI + Pb(NO
3
)
2
→ KNO
3
+ PbI
2
Slide 8
Types of Chemical Reactions
1.
Combination or Synthesis
2 or more reactants → 1 product
Example: H
2
+ O
2
→ H
2
O
2. Decomposition
1 reactant → 2 or more products
Example: H
2
O
2
→ H
2
O + O
2
Slide 11
Types of Chemical Reactions
Classify each reaction:
1.
NaCl + AgNO
3
→ AgCl + NaNO
3
2.
Na + H
2
O → NaOH + H
2
3.
4.
5.
C
2
H
6
+ O
2
→ CO
2
+ H
2
O
Mg + O
2
→ MgO
Na
2
CO
3
→ Na
2
O + CO
2
6.
KOH + HCl → KCl + H
2
O

Slide 10
Types of Chemical Reactions
5.
Combustion fuel + oxygen → water + carbon dioxide
Example: CH
4
+ O
2
→ CO
2
+ H
2
O

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) CaCO
3
→ CaO + CO
2
(b) BaCl
2
+ Na
2
SO
4
→ BaSO
4
+ 2NaCl
(c) 3HNO
3
+ Al(OH)
3
→ 3H
2
O + Al(NO
3
)
3
(d) 2C
2
H
2
+ 5O
2
→ 4CO
2
+ 2H
2
O
(e) Na
2
O + H
2
O → 2NaOH
____________________________________
____________________________________
____________________________________
____________________________________
____________________________________
(f) Mg + 2HCl → MgCl
2
+ H
2
____________________________________
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 (H
2
).
What are the reactants? ________________________________________________________
What are the products? ________________________________________________________
Chemical Equation: ___________________________________________________________
B.) Iron reacts with oxygen (O
2
) 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) AgNO
3
( aq ) + NaCl( aq
) →
C) Mg( s ) + CuSO
4
( aq
) →
D) Ag
2
O →
E) KOH( aq ) + BaCl( aq
) →
F) C
4
H
10
( g ) + O
2
( g ) →
G) Na( s ) + O
2
( 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:
2
2
2
2
2
2
2
2
2
2
6
3
12
6
2
2

Solve the following stoichiometric problems. Show your work.
8. Mole-Mole Problems
In the chemical reaction, Mg + 2HCl → MgCl
2
+
A. how many moles of magnesium are needed to produce 3.00 moles of hydrogen gas?
H
2
B. how many moles of magnesium chloride can be produced from 4.00 moles of hydrogen chloride?
In the reaction, 4Fe + 3O
2
→ 2Fe
2
O
3
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 → MgCl
2
A. how many moles of magnesium are needed to produce 4.00g of hydrogen gas?
+ H
2
B. how many grams of magnesium chloride can be produced from 2.00 moles of magnesium?
In the reaction, 4Fe + 3O
2
→ 2Fe
2
O
3
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, C
6
H
12
O
6
(g) + 6O
2
(g) → 6CO
2
(g) + 6H
2
O(g)
A. how many grams of carbon dioxide can be produced from the burning of 180.0 g of glucose
(C
6
H
12
O
6
)?
B. what is the mass of oxygen needed to produce 54.0 g of water?

6
20
10
In the chemical reaction, 2Mg + CO
2
→ 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) → MgCl
2
(aq)
A) how many liters of hydrogen gas at STP is produced from 2.50 moles of magnesium?
+
B) how many liters of hydrogen gas at STP is produced from 18.0 g of magnesium?
In the chemical reaction, C
6
H
12
O
6
(g) + 6O
2
(g) → 6CO
2
(g) + 6H
2
O(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 → 2 sandwiches
For the burger “reaction”, complete the table below:
Number of
Bread Slices
Number of Ham
Slices
Number of
Sandwiches
Name of Excess
“Reactant”
10 20
Name of
Limiting
“Reactant”
H
2
(g)
Excess amount
20 24
6
12
12 ham bread
20 slices
10 slices

B) Hydrogen gas reacts with oxygen gas to form water vapor according to the reaction below:
2H
2
( g ) + O
2
( g ) → 2H
2
O( g )
For this reaction, complete the table below:
Amount of
Hydrogen
Amount of
Oxygen
Amount of
Water Vapor
Name of
Limiting
Reactant
Name of Excess
Reactant
10 moles 8 moles
Excess Amount
20 moles
4 g
6 moles
40 g
10 g
2 L
12 L
64 g
2 L
4 L
13. Theoretical and Percent Yield
A. In the chemical reaction, 2Mg + CO
2
→ 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?
N
2
+ 3H
2
→ 2NH
3

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 + O
2
→ Al
2
O
3
1. What is/are the reactant(s) of the reaction shown above?
2. What type of reaction is it?
Explain.
Direction Instruction:
Predicting Products
Guided Practice
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.
Students look at other chemical reactions. Working in groups, they practice identifying reactants and products and writing chemical equations.
Youtube video of reaction between sodium and chlorine – see link below
Chemical Reactions and Stoichiometry
PowerPoint- see
Supplementary
Materials folder
10 minutes
30 minutes
Check for
Understanding/Exit
Ticket
Provide evidence of mastery of the day’s learning objectives.
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.
Homework Practice Worksheet
Question 4
Youtube video of reaction between sodium and chlorine http://www.youtube.com/watch?v=Mx5JJWI2aaw
See below -Chemical
Reactions and
Stoichiometry
Practice Worksheet

PowerPoint Slide Master
Slide 12
Learning Objective:
 Write chemical equations from word equations.
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 (O
2 air and burns to produce magnesium oxide (MgO).
) in
Reactants: magnesium (Mg) and oxygen (O
2
)
Products: magnesium oxide (MgO)
Chemical Equation: Mg + O
2
 MgO
Slide 15
Slide 16
Write the chemical equation for this reaction:
Blue copper(II) sulfate solution reacts with iron to form iron(II) sulfate and copper.
Reactants:
Copper(II)sulfate and iron
Products:
Iron (II)sulfate and copper
Chemical Equation:
CuSO
4
+ Fe → FeSO
4
+ Cu
Write the chemical equation for each reaction:
1.
Silver oxide decomposes into silver and oxygen gas when heated.
2.
Ethanol (C
2
H
5
OH) 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)
2MCuCl
2
, Al
1MBaCl
2
, 1M Na
2
SO
4
Demo double displacement reactions
Differentiate between single and double displacements
1M BaCl
2
,
1M Na
2
SO
4
Time Notes
5 minutes Observe reaction, understand mechanism of different types of reactions :
5 minutes
1. Observe reaction,
2.
write balanced equation,
3.
Predict products and write balanced equations
Introduce synthesis,decomposition
& combustion
Check for understanding
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
Objective:
 Predict the products of common chemical reactions.
Slide 18
Predicting Products
Al + HCl → _________
 What type of reaction will most likely occur between the 2 reactants?
 What are the products?
Al + HCl → AlCl
3
+ H
2
Slide 19
Predicting Products
CuCl
2
( aq ) + Pb(NO
3
)
2
( aq
) → _________
 What type of reaction will most likely occur between the 2 reactants?
 What are the products?
CuCl
2
( aq ) + Pb(NO
3
)
2
( aq ) → Cu(NO
3
)
2
( aq ) + PbCl
2
( s )
Slide 20
Predicting Products
C
6
H
14
+ O
2
→ _________
 What type of reaction will most likely occur between the 2 reactants?
 What are the products?
C
6
H
14
+ O
2
→ CO
2
+ H
2
O
Slide 21
Predict the products of the reaction:
1.
NaOH( aq ) + FeCl
3
( aq
) →
2.
3.
4.
5.
Zn( s ) + HNO
3
( aq ) →
C
4
H
10
( g ) + O
2
( g
) →
N
2
( g ) + H
2
( g ) →
KBr( aq ) + Cl
2
( 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 Objective Materials Required Time
Warm Up Students predict whether mass will increase, decrease or decrease in the reaction between
Alkaseltzer and water in a sealed ziplock bag.
They will also write a simple procedure for testing their prediction.
Alkaseltzer Mini-Lab Students test their prediction.
Per group:
1 ziplock bag
1 Alkaseltzer tablet cup with water triple beam balance
10 minutes
10 minutes
10 minutes
Notes
To give students hints on how to write the experimental procedure, provide them a list of materials they may use.
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.
Direct Instruction: Law of Conservation of
Mass and Balancing
Equations
Guided Practice on
Balancing Equations
Chemical Reactions and
Stoichiometry
PowerPoint- see
Supplementary
Materials folder
20 minutes

Check for
Understanding/Exit
Ticket
Homework
1) 24 grams of carbon completely reacts with
64 grams of oxygen gas.
What mass of carbon dioxide is produced?
C + O
2
 CO
2
24g 64g ?
2) Balance the following equation:
Al + Fe
2
O
3
→ Al
2
O
3
+ Fe
Practice Worksheet,
Questions 6-7
See attached
Chemical Reactions and Stoichiometry
Practice Worksheet

PowerPoint Slide Master
Slide 22
Objectives:
 Recognize that chemical reactions are governed by the Law of Conservation of Mass.
 Balance chemical equations.
Slide 26
Balanced Chemical Equations
Balanced Equation – the number of atoms of each element is equal on both sides of the equation
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.
Slide 23
Law of Conservation of Mass
Burning Magnesium Metal in an Open Container
Slide 27
Balancing Chemical Equations
Example 1:
Mg +
Reactants:
2 HCl →
Mg – 1
H – 1
X 2 = 2
Cl – 1
X 2 = 2
MgCl
2
Products:
Mg – 1
H – 2
Cl - 2
+ H
2
Slide 24
Edition, 1990, page 77
Law of Conservation of Mass
Burning Magnesium Metal in a Closed Container
Slide 28
Balancing Chemical Equations
Example 2:
2 Na + 2 H
Reactants:
2
O → 2 NaOH+ H
2
Products:
Na – 1
X 2 = 2
H – 2
X 2 = 4
O – 1
X 2 = 2
Na – 1
X 2 = 2
O - 1
X 2 = 2
= 4
Slide 25
Edition, 1990, page 77
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 + O
2
→ 2MgO
48 g 32 g ?
Slide 30
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
2.
Fe
3.
Zn
+
+
+
Cl
2
O
2
HCl
4.
KNO
3
→ KNO
2
→ NaCl
→ Fe
2
O
3
→ ZnCl
2
+ H
2
+ O
2

Day 13
(Essential Question(s): ____________________________________________________________________________________)
Learning Objectives:

Balance chemical equations.
Standard 3a
Activity
Warm Up
Independent Practice
– Balancing Equations
Objective
Review balancing of equations.
This will allow students to practice further and master the skill of balancing equations
Materials Required
See Pogil-Balancing
Equations (pdf file)
Time
5 minutes
35 minutes
Notes
KClO
3
→ KCl + O
2
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.
Check for
Understanding/Exit
Ticket
Students write a balanced chemical equation for the reaction between iron and sulfur.
You may show a video of the reaction or demonstrate it yourself.
Youtube video of the reaction – see link below.
Demo – see Activity
Sheet below
10 minutes
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
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
Materials Required
Rotational Lab Stations
– Types of Reactions
Check for
Understanding/Exit
Ticket
Homework
Objective
Go over the objective, procedure and safety precautions of the lab.
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
See above - Chemical
Reactions and
Stoichiometry Practice
Worksheet
Time
10 minutes
40 minutes
5 minutes
Notes

Purpose: Predict products of a reaction and write chemical equations.
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: _________________________________________________________
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: _________________________________________________________

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: _________________________________________________________
2
5
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: _________________________________________________________

4
∙
2
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
Direct Instruction:
Mole-Mole Conversion
Guided Practice
Check for
Understanding/Exit
Ticket
Homework
Objective
Students interpret a recipe and relate it to interpreting balanced chemical equations.
Show students how the conversion is done through dimensional analysis.
Provide evidence of mastery of the day’s learning objective.
Practice Worksheet,
Question 8, 1-d
Materials Required
Chemical Reactions and
Stoichiometry
PowerPoint- see
Supplementary
Materials folder
See above - Chemical
Reactions and
Stoichiometry Practice
Worksheet
Time Notes
10 minutes 1) 1 bun + 2 patties + 2 cheese
→ 1 cheeseburger
How many patties are needed to make 5 cheeseburgers?
2) 2H
2
+ O
2
→2H
2
O
How many moles of hydrogen are needed to make 4 moles of water?
15 minutes Emphasize that mole ratio
(conversion factor) is based on
balanced equation. The
coefficients indicate number of moles.
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.
10 minutes
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 N
2
to produce 2.50 moles of NH
3
.
N
2
+ 3H
2
→ 2NH
3
Support your answer with a calculation.

PowerPoint Slide Master
Slide 31
Learning Objective:
 Convert moles of reactants to moles of products and vice versa.
Slide 32
Stoichiometry
-Stoichiometry is the calculation of the amount of reactants and products in a chemical reaction.
Amount is usually expressed in number of moles , mass or volume (gases).
Stoichiometric calculations are based on balanced equations .
Slide 35
Mole-Mole Conversion
2 H
2
+ O
2
→ 2 H
2
O
Ex.1 : How many moles of water can be produced from 3.50 moles of hydrogen?
Given: 3.50 mol H
2
Unknown: mol H
2
O
Possible conversion factors: 2 mol H
2
2 mol H
2
O
Derived from balanced equation
2 mol H
2
O
2 mol H
2
Equation: moles of H
2
O = 3.50 mol H
2 x 2 mol H
2
O = 3.50 mol H
2
2 mol of H
2
O
Slide 36
Mole-Mole Conversion
2 H
2
+ O
2
→ 2 H
2
O
Ex. 2: How many moles of hydrogen are needed to react with 5.0 moles of oxygen?
Given: 5.0 mol O
2
Possible conversion factors:
Equation: moles of H
2
= 5.0 mol O
2
2mol H
2
1 mol O
2
1 mol O
2
2mol H
2
Unknown: mol H
2 x 2 mol H
2
1 mol of O
2
=
10. moles H
2
Slide 37
Practice Problems:
Mg + 2 HCl → MgCl
2
+ H
2
1. How many moles of magnesium are needed to produce
0.500 moles of magnesium chloride?
2. How many moles of hydrogen gas can be produced from 6 moles of magnesium?
Slide 33
Solving Stoichiometric Problems
1.Identify given and unknown.
2.Write possible conversion factors.
3. Set up equation using appropriate conversion factor(s). mole unknown = mole of known x mole of unknown mole of known
Do the known units cancel?
4. Check answer. Sig figs? Units?
Slide 34
Mole-Mole Conversion
2 H
2
2 moles
+ O
2
1 mole
→ 2 H
2
O
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

Day 18
(Essential Question(s): ____________________________________________________________________________________)
Learning Objective:

Do mass-mole conversions of reactants and products.
Standard 3a, e
Activity
Warm Up
Direction Instruction –
Mole-Mass Conversion
Guided Practice
Objective
This is a review of molar mass to prepare students to do mass-mole conversions.
Show students how the conversion is done through dimensional analysis.
Materials Required
Chemical Reactions and
Stoichiometry
PowerPoint- see
Supplementary
Materials folder
Time
10 minutes
15 minutes
20 minutes
Notes
What is the mass of a mole of:
1) N
2
H
4
2) Ca(NO
3
)
2
Emphasize: Molar mass is used to do mole-mass conversion.
Check for
Understanding/Exit
Ticket
Homework
Provide evidence of mastery of the day’s learning objective.
Practice Worksheet,
Question 9, 1-d
See above - Chemical
Reactions and
Stoichiometry Practice
Worksheet
10 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?)
2KClO
3
→ 2KCl + 3O
2
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.

PowerPoint Slide Master
Slide 38
Mole-Mass or Mass-Mole Conversion
Given Molar mass
Mole ratio from balanced equation mol B = mass A x 1 mol A x mol B mass A mol A mass B = mol A x mol B x mass B mol A 1 mol B
Slide 39
Slide 40
Mole-Mass or Mass-Mole Conversion
2H
2
+ O
2
→ 2H
2
O
Ex.1 :
What is the mass of oxygen that is needed to produce 4.0 moles of water?
Given: 4.0 mol H
2
O Unknown: g of O
2
Possible conversion factors: 1 mol O
2
32g O
2
2 mol H
2
O 1 mol O
2
Equation: mass of O
2
= 4.0 mol H
2
O x 1 mol O
2 mol H
2 x 32g O
2
2
O 1 mol O
2
= 64 g O
2
Practice Problems:
Mg + 2 HCl → MgCl
2
+ H
2
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
Direct Instruction –
Mass-Mass Conversion
Guided Practice
Objective
Review mole-mass conversion.
Show students how the conversion is done through dimensional analysis.
Materials Required
Chemical Reactions and Stoichiometry
PowerPoint- see
Supplementary
Materials folder
Time
10 minutes
15 minutes
20 minutes
Notes
N
2
+ 3H
2
→ 2NH
3
How many grams of ammonia can be produced from 1.50 moles of hydrogen gas?
Check for
Understanding/Exit
Ticket
Homework
Provide evidence of mastery of the day’s learning objectives.
Practice Worksheet,
Question 10, 1-d
10 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?)
Mg + CuSO
4
→ MgSO
4
+ Cu
How many grams of Cu can be produced when 5.00 g of
Mg reacts completely with
CuSO
4
?
(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).
See above - Chemical
Reactions and
Stoichiometry Practice
Worksheet

PowerPoint Slide Master
Slide 41
Mass- Mass Conversion
2 H
2
2 (2g) = 4g
+ O
2
1 (32g) = 32 g
→
2 H
2
O
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 x mole B x mass B mass A mole A 1 mole B
2H
2
+ O
2
→ 2H
2
O
Ex.1 : What is the mass of oxygen that is needed to produce 18 g of water?
Given: 18 g of H
2
O Unknown: g of O
2
Possible conversion factors: 1 mol H
2
O 1 mol O
2
32g O
2
18 g H
2
O 2mol H
2
O 1 mol O
2
Equation: mass of O
2
= 18 g H
2
O x 1 mol H
2
O x 1 mol O
2 x 32 g O
2
18 g H
2
O 2 mol H
2
O 1 mol O
2
= 16 g O
2
Slide 43
Practice Problems:
Mg + 2 HCl → MgCl
2
+ H
2
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
Demo – Hydrogen
Balloon Explosion
Guided Practice
Objective
Review mass-mole conversion.
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 H
2 and O
2
, let the students figure out the volume of O
2 needed to completely burn the H
2
gas.
Materials Required
See Activity Sheet below
Time
5 minutes
25 minutes
20 minutes
Notes
Zn + HCl → ZnCl
2
+ H
2
How many moles of hydrogen gas can be prepared from the reaction of 2g of Zn with excess HCl?
Be aware of the safety precautions that must be observed in demonstrating the ignition of H
2
gas. See
Activity Sheet below.
Chemical Reactions and
Stoichiometry
PowerPoint- see
Supplementary
Materials folder
5 minutes Check for
Understanding/Exit
Ticket
Homework
2C
4
H
10
+ 13O
2
→ 8CO
2
+
10H
2
O
If 0.33 moles of butane
(C
4
H
10
) are burned, how many liters of carbon dioxide would be produced at STP?
Practice Worksheet,
Question 11 a-d
See above - Chemical
Reactions and
Stoichiometry Practice
Worksheet

PowerPoint Slide Master
Slide 44
Learning Objectives:
 Perform mass-mole-volume conversion at STP.
 Perform volume-volume conversion of gaseous reactants and products at STP.
Slide 45
Molar Volume
 Avogadro’s Principle: Equal volumes of gases at the same temperature and pressure contain the same number of particles.
 At STP (Standard Temperature and Pressure),
1 mole of any gas occupies a volume of 22.4 L .
Slide 46
Mass-Mole-Volume Conversion (at STP)
L of B = mass of A x 1 mol A x mol B x 22.4 L B mass A mol A 1 mol B
2 KClO
3
→ 2KCl + 3O
2
Ex.1 : How many liters of oxygen gas are produced when 30.0 g of potassium chlorate decomposes at STP?
Given: 30.0 g KClO
3
Unknown: L of O
2
Possible conversion factors: 1 mol KClO
3
122.5 g KClO
3
Equation:
L of O
2
3mol O
= 30.0 g KClO
3 x 1 mol KClO
3
122.5 g KClO x 3 mol O
2
3
2 mol KClO
3
2
2mol KClO
3
22.4 L O
1 mol O x 22.4 L O
2 mol O
2
2
2
= 8.23 L O
2
Slide 47
Practice Problems:
Mg + 2 HCl → MgCl
2
+ H
2
1. How many grams of magnesium are needed to produce 11.2
L of hydrogen gas at STP?
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?
Slide 48
Slide 49
Volume -Volume Conversion (at STP)
L of B = L of A x mol B mol A
2H
2
+ O
2
→
2H
2
O
Ex.1 : How many liters of oxygen gas are needed to completely react with 13.5 L of hydrogen gas at STP?
Given: 13.5 L H
2
Possible conversion factors: 1mol O
2
2 mol H
2
Equation:
Unknown: L of O
2
L of O
2
=
13.5 L H
2 x
1mol O
2 = 6.75 L O 2
2 mol H
2
Practice Problems:
N
2
(g) + 3H
2
(g) → 2 NH
3
(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
Materials Required Time
10 minutes
Objective
1 bun + 2 patties + 2 cheese slices → 1 cheeseburger
If there are 8 buns, 12 patties and 12 cheese slices available, how many cheeseburgers can we make?
15 minutes
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.
Direction Instruction –
Solving limiting and excess reactant problems
Guided Practice
Check for
Understanding/Exit
Ticket
Homework
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
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?
2H
2
+ O
2
→ 2H
2
O
3. If 1.50 moles H
2
and
0.50 moles O
2
react, will both reactants be completely consumed?
If not, name the excess reactant.
Practice Worksheet –
Question 12 a-b
Chemical Reactions and Stoichiometry
PowerPoint- see
Supplementary
Materials folder
See above - Chemical
Reactions and
Stoichiometry Practice
Worksheet
20 minutes
10 minutes

PowerPoint Slide Master
Slide 50
Learning Objective:
 Distinguish between limiting and excess reactants.
Slide 51
Limiting and Excess Reactants
1 bun + 2 patties + 2 cheese slices → double cheeseburger
If there are 5 buns, 8 patties and 6 cheese slices available, how many double cheeseburgers can be made?
Which ingredient is completely used up?
Which ingredient is left over?
Slide 52
Limiting and Excess Reactants
Limiting Reactant – completely used up; limits the amount of product
Excess Reactant
– not completely used up, “left over”
Slide 53
Slide 54
Limiting and Excess Reactants
2H
2
+ O
2
→
Ex.1: 6.0 g of H
2 and 60.g of O
2 are made to react.
2H
2
O
(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 O
2
= 6.0 g H
2 x 1 mol H
2
2 g H
2 x 1 mol O
2 x 32g O
2
= 48 g O
2 mol H
2
1 mol O
2
2
Only 48 g of O
2
O
2 is needed to completely react, so it is an excess reactant. 12 g of is left over.
Limiting and Excess Reactants
2H
2
+ O
2
→
Ex.1: 6.0 g of H
2 and 60.g of O
2 are made to react.
2H
2
O
(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?
H
2 is the limiting reactant; it determines the amount of water produced.
mass of H
2
O = 6.0 g H
2 x 1 mol H
2
2 g H
2 x 2 mol H
2
O x 18 g H
2
O = 54g H
2
O
2 mol H
2
1 mol H
2
O
Slide 55
Practice Problems:
2Al + 3Br
2
→ 2AlBr
3
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 Description/Details
Warm Up Review mass-mass conversion.
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
Pre-Lab Discussion –
Preparation of Salt
Check for
Understanding/Exit
Ticket
Homework
Go over objective, procedure and safety precautions of the lab.
Provide evidence of mastery of the day’s learning objective.
Answer Pre-lab
Questions –
Preparation of Salt
Materials Required
Chemical Reactions and Stoichiometry
PowerPoint- see
Supplementary
Materials folder
Time
10 minutes
30 minutes
10 minutes
Notes
2Na + Cl
2
→ 2NaCl
What is the maximum amount (in grams) of NaCl that can be produced if 56.0 g sodium reacts with excess chlorine?
Chemists often get less than the maximum amount of product that they expect from a reaction. Why do you think this is so?
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.
NaHCO
3
+ HCl → NaCl + H
2
O + CO
2
How many grams of NaCl may be produced if 8.40g NaHCO
3
reacts completely with excess HCl solution?
What is the percent yield of the reaction, if only 4.0 g NaCl is actually produced?

PowerPoint Slide Master
Slide 56
Learning Objective:
 Determine the theoretical and percent yield of a reaction.
Slide 59
Slide 60
Sample Problem 2
Mg + 2 HCl → MgCl
2
+ H
2
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: % yield = 1g x 100 = 50%
2g
Check for Understanding
2KClO
3
→ 2KCl + 3O
2
What is the % yield of the above reaction if only 45 g of oxygen is produced from 122 g of potassium chlorate?
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 = Actual
Theoretical x 100
Slide 58
Sample Problem 1
2H
2
2(2g) = 4g
+ O
2
1(32g) = 32 g
→ 2H
2
O
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

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 Materials Required
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.
See Activity Sheet below for list of materials
Answer analysis questions and write conclusion.
Time
5 minutes
45 minutes
5 minutes
Notes

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
(NaHCO
3
), commonly known as baking soda, and hydrochloric acid (HCl). The reaction produces table salt or sodium chloride (NaCl) and carbonic acid (H
2
CO
3
). The carbonic acid readily decomposes to water and carbon dioxide as shown by the equations below.
NaHCO
3
+ HCl → NaCl + H
2
CO
3
NaHCO
3
+ HCl → NaCl + H
2
O + CO
2
A known mass of NaHCO
3
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 (Ag
2
O) was heated several times to drive out all of the oxygen, leaving behind silver, according to the reaction,
2Ag
2
O → 4Ag + O
2
The following data were collected.
Mass of crucible (g)
Mass of crucible + Ag
2
O (g)
Mass of crucible + remaining solid after first heating (g)
Mass of beaker + remaining solid after second heating (g)
Mass of beaker + remaining solid after third heating (g)
20.5552
22.5535
22.38220
22.1621
22.1621
(1) How many grams of silver oxide (Ag
2
O) 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 (NaHCO
3
) and hydrochloric acid (HCl).
Procedure:
1.
Weigh an empty beaker.
2.
Place about 1 g of NaHCO
3
into the beaker. Record the mass of the beaker and the NaHCO
3
.
3.
Add HCl to NaHCO
3
drop by drop. When the reaction mixture stops fizzing all of the NaHCO
3 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 + NaHCO
3
(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 NaHCO
3
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 NaHCO
3
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

Slide 1
Slide 2
Slide 3
Slide 4
Slide 5
Test Review
Chemical Reactions and
Stoichiometry
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.
CaCO
3
→ CaO + CO
2
Which is true of the above reaction?
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.
B
Which is a single displacement reaction?
A. NH
4
NO
2
B. Na
2
SO
4
→ N
2
+ 2H
+ Ba(NO
3
)
C. Fe + 2HCl → FeCl
2
2
2
O
→ BaSO
+ H
2
D. C
2
H
5
OH + 3O
2
→ 2CO
2
4
+ 2NaNO
+ 3H
2
O
3
C
Which of these equations is balanced?
A. NH
4
B. NH
3
NO
2
→ N
2
+ H
+ HCl → NH
4
2
Cl
O
C. 2NH
3
→ N
2
D. NH
4
NO
3
+ H
→ N
2
2
O + H
2
O
B
Slide 7
Slide 8
Slide 9
Slide
10
Slide
11
NH
3
+ HCl → NH
4
Cl
How many grams of NH
3 are needed to produce 25 g of ammonium chloride?
8.0 g NH
3
2NH
3
→ N
2
+ 3H
2
How many moles of NH
3 are needed to produce 2.0 g of hydrogen?
0.67 moles NH
3
Propane reacts with oxygen to produce water and carbon dioxide based on the equation:
C
3
H
8
+ 5O
2
→ 4H
2
O + 3CO
2
What volume of propane at STP is needed to produce 6.0 moles of CO
2
?
44.8 L
Propane reacts with oxygen to produce water and carbon dioxide based on the equation:
C
3
H
8
+ 5O
2
→ 4H
2
O + 3CO
2
How many grams of oxygen gas are needed to completely burn 22.4 L of propane at STP?
160 g
2H
2
O → 2H
2
+ O
2
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 O
2
, 94%

Slide 6
NH
3
+ HCl → NH
4
Cl
How many moles of hydrogen chloride are needed to produce
0.35 moles of ammonium chloride?
0.35 moles HCl
Slide
12 4Al+ 3O
2
→ 2Al
2
O
3
100.0 g of Al and 100.0 g of O
2 react. What are the limiting and excess reactants? How many grams of Al
2
O
3 are produced?
Limiting: Al Excess: O
2
188.9 g Al
2
O3

Slide 1
Slide 2
Slide 3
Slide 4
Slide 5
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
“Matter is made up of atoms. Atoms have tiny positive centers containing protons and neutrons.”
The above statement is a/an
A. inference
B. hypothesis
C. theory
D. observation
Answer: C
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
Answer: D
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 7
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 8
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 9
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
Slide 10
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
Slide 12
The nucleus of an atom
A. is negatively-charged
B.
accounts for most of the atom’s mass
C.
occupies most of the atom’s volume
D. contains electrons and protons
Answer: B
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
Give the number of protons, electrons and neutrons of 15
8
O.
Answer: electrons- 8, protons- 8, neutron-7
Slide 17
How many valence electrons does nitrogen have?
Answer: 5
Slide 13
Slide 14
Slide 15
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: A
Isotopes of the same element always have the same
A. atomic number
B. mass number
C. atomic mass
D. number of neutrons
Answer: A
Slide 18
Draw the Lewis Dot diagram of carbon
Answer:
C
Slide 19
Which element do you expect to have the same valence electrons as carbon?
A. calcium
B. iron
C. chlorine
D. silicon
Answer: D
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: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1

Slide 21
Which of these is a halogen?
A. helium
B. carbon
C. lithium
D. iodine
Answer: D
Slide 26
Which element is the least electronegative?
A. Nitrogen
B. Phosphorus
C. Oxygen
D. Sulfur
Answer: B
Slide 22
Slide 23
Slide 24
Slide 25
To which family does the element krypton belong?
A. Alkali metal
B. Alkaline-earth metal
C. Transition Metal
D. Noble gases
Answer: D
Slide 27
Which type of chemical bond exists between a metal and a non-metal?
Ionic
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
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 28
Answer: A
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 (C
2
H
5
OH) solution
C. Potassium chloride (KCl) solution
D. Hydrochloric acid (HCl) solution
Answer: B

Slide 31
Slide 32
Slide 33
Slide 34
Slide 35
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
Slide 36
Give the name of
FeSO
4
Iron (II) sulfate
What is the correct formula of sodium bromide?
A. NaBr
B. Na
2
Br
C. NaBr
2
D. Na
2
Br
2
Answer: A
Slide 37
What is the correct formula of sodium bromide?
A. NaBr
B. Na
2
Br
C. NaBr
2
D. Na
2
Br
2
Answer: A
Slide 38
What is the chemical formula for sodium sulfate?
Na
2
SO
4
Which of these is an acid?
H
2
SO
4
KOH
K
2
SO
4
Answer: H
2
SO
4
Slide 39
What is the chemical formula for copper (II) nitrate?
Cu(NO
3
)
2
Which is a polar covalent bond?
A. N-H
B. As-H
C. O-O
D. Cl-Cl
A
Slide 40
Give the name of
CaCl
2
Calcium Chloride
Write the Lewis dot diagram of
H
2
S.
H S H

Slide 41
Slide 42
Slide 43
Slide 44
Slide 45
Slide 46
How many non-bonding (unshared) pairs of electrons does sulfur have in
H
2
S?
2
Predict the shape of the NH
3
.
Trigonal Pyramidal
How many non-bonding pairs does sulfur have in H
2
S?
2
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
C,D,A,B
What is the shape of the H
2
S molecule?
Bent
Slide 48
Which of these is an alkene?
A. C
3
H
8
B. C
4
H
8
C. C
5
H
8
D. C
6
H
14
B
Slide 49
Is the H
2
S molecule polar or nonpolar?
Polar
B
Which of these is a polar molecule?
A. CO
2
B. SF
6
C. CH
4
D. PCl
3
D
Slide 50
NH
3
+ HCl → NH
4
Cl
How many grams of NH
3 produce 25 g of NH
4
Cl?
is needed to
8.0 g of NH
3

Day 28, 29, 30