Advanced Chemistry
Student Packet - All Units of Study (2023-2024)
Lesson 1 - Course Introduction
Unit 1 - Moles and Masses - Learning Goal Summary
Lesson 2 - The Mole, Mass and Avogadro’s Number
Lesson 3 - Empirical and Molecular Formulas
Lesson 4 - Balancing Reactions Review and Hydrates
Lesson 5 - Formative Lab #1: Empirical Formulas
Lesson 6 - Formative Quiz 1
Introduction to Mass-Mass and # of Particle Relationships and Limiting Reagents
Lesson 7 - Limiting and Excess Problems, Actual and Percent Yield
Lesson 8 - Formative Lab #2: Limiting Reagents
Lesson 9 - Unit 1 Summative Lab - Not in this Packet
Lesson 10 - Formative Quiz 2 Combined/Unit Review
Lesson 11 - Unit 1 Moles and Masses Summative Test - Not in this Packet
Unit 2 - Gases and Solutions - Learning Goal Summary
Lesson 12 - Unit 2 Lesson 1 - Introduction to Ideal Gases and KMT + Demos
Lesson 13 - Unit 2 Lesson 2 - Combined Gas Law + Demos
Lesson 14 - Unit 2 Lesson 3 - Ideal Gas Law
Lesson 15 - Unit 2 Lesson 4 - Deviations from Ideal Gas Behaviour
Lesson 16 - Unit 2 Lesson 5 - Formative Quiz 1/Gas Stoichiometry and Limiting Reagents
Lesson 17 - Unit 2 Lesson 6 - Formative Lab 1 - Molar Mass of a Gas
Lesson 18 - Unit 2 Lesson 7 - Introduction to Solutions and Concentration
Lesson 19 - Unit 2 Lesson 8 - Making Standard Solutions and Dilutions
Lesson 20 - Unit 2 Lesson 9 - Beer-Lambert Law Theory and Demo
Lesson 21 - Unit 2 Lesson 10 - Intro to Titrations
Lesson 22 - Unit 2 Lesson 11 - Formative Lab 2 - Standardization and Titration
Lesson 23 - Unit 2 Lesson 12 - Summative Lab - Not in this Packet
Lesson 24 - Unit 2 Lesson 13 - Formative Quiz 2/Unit 2 Review
Lesson 25 - Unit 2 Lesson 14 - Unit 2 Gases and Solutions Summative Test - Not in Packet
Unit 3 - Bonding and Electrons - Learning Goal Summary
Lesson 26 - Unit 3 Lesson 1 - Lewis Structures and Formal Charges
Lesson 27 - Unit 3 Lesson 2 - VSEPR Theory, 1-4 Domains (shapes and bond angles)
Lesson 28 - Unit 3 Lesson 3 - Review of IMFs, Determination of Polarity
Lesson 29 - Unit 3 Lesson 4 - Formative Quiz 1/Chromatography
Lesson 30 - Unit 3 Lesson 5 - Formative Lab 1 - Polarity and Chromatography
Lesson 31 - Unit 3 Lesson 6 - Cumulative Design Lab - Day 1
Lesson 32 - Unit 3 Lesson 7 - Intro to Redox Processes
Lesson 33 - Unit 3 Lesson 8 - Balancing Redox Reactions
Lesson 34 - Unit 3 Lesson 9 - Voltaic Cells
Lesson 35 - Unit 3 Lesson 10 - Formative Quiz 2/Unit Review
Lesson 36 - Unit 3 Lesson 11 - Advanced Chemistry Semester Exam
Lesson 37 - Unit 3 Lesson 12 - Cumulative Design Lab - Day 2
Lesson 38 - Unit 3 Lesson 13 - Cumulative Design Lab - Day 3
Lesson 39 - Unit 3 Lesson 14 - Exam Corrections and Lab Peer-Assessment
Lesson 40 - TBD
1
Lesson 1: Advanced Chemistry Course Overview
Welcome to Advanced Chemistry!
Here are the units of study for this year:
Unit 1 - Moles and Masses - 11 Lessons
Unit 2 - Gases and Solutions - 14 Lessons
Unit 3 - Bonding and Electrons - 14 Lessons (approximately)
How to be successful in Advanced Chemistry?
1. Complete Homework
Complete all homework assignments on time. Almost all assignments are due before the following
lesson.
2. Check your Answers to Homework Exercises (all students should be doing this for every
assignment). Then, refer to your teacher for challenging concepts.
3. Ask Lots of Questions
● You can ask questions to me in class or to your classmates. At the end of every class I
invite students who still have questions to stay behind to ask them.
● Use google calendar to set up appointments with me during your honor pass or during my
community day office hours.
4. Use the Learning Goals as a Checklist
We share learning goals with you for every lesson, they are in the notes and on the Advacned
Chemistry Online Learning Doc. Read through these learning goals and use them as a checklist
as you prepare for formative quizzes and summative tests.
5. Study for Tests Well in Advance
This is an example of distributed practice. Prepare a schedule of short study sessions for
chemistry throughout the unit rather than trying to cram everything in the night before the test.
6. Apply Effective Learning Techniques When You’re Studying
Spaced Practice
Retrieval Practice (quizlet.com, worksheets, formative quiz V2, are great ways to practice retrieval)
Elaboration
Interleaving
2
Concrete Examples
Dual Coding
How will your final letter grade be determined in Advanced
Chemistry?
JIS’ assessment policy will apply to this course. Therefore, your letter grade will be determined at
the end of the course based on an evaluation of your ‘most recent and consistent’ learning level.
You can come and see your teacher anytime if you have questions about your standing in the
course.
Here are the grade breakdowns for Advanced Chemistry.
Advanced Chemistry Grade Breakdown
Summative Tests and Summative Labs
85%
Final Exam
15%
Advanced Chemistry Grade Scale (HS4)
Grade
Cutoff
A*
90%
A
80%
B
65%
C
50%
D
35%
F
< 35%
3
Unit 1: Moles and Masses - Learning Goal Summary
Unit 1: Moles and Mass “Can Do Review”
Can You? / Do You?
UNDERSTANDING
Can you define the term mole as a sample of particles, atoms, molecules, ions, etc. with
6.02 x 1023 species?
Do you know all of the polyatomic ions?
Do you know the names and formulae of common compounds—water, ammonia,
hydrochloric acid, sulphuric acid, nitric acid, etc.?
Do you know the oxidation states of the main elements?
Do you know oxidation states for common transition elements—Ag+, Zn2+, Pb2+/Pb4+
Do you know the difference between theoretical and experimental yield?
EXPLAINING
Can you write names and formulae for ionic and covalent compounds?
Can you write symbolic equations from word equations?
Can you write word equations from symbolic equations?
Can you balance equations?
PROBLEM SOLVING
Can you determine the molar mass of a given compound?
Can you use molar mass and/or Avogadro’s number to convert mass, moles and number of
particles?
Can you determine the formula of a hydrate from experimental data?
Can you determine empirical and molecular formulae using experimental data?
Can you solve problems involving chemical equations, mass and number of chemical
species?
Can you create ratios between the coefficients in balanced equations to solve unknown
values?
Can you calculate percent yield based on experimental yield and calculated theoretical
yield?
Can you identify the limiting and excess reagent in a problem?
Can you solve problems involving limiting reagents?
Can you calculate the moles, mass or number of species in excess based on the amount of
limiting reagent used?
EXPERIMENTAL SKILLS
Can you create and/or use a correct data table? This includes correct labels, units and
consistent use of decimal places.
Can you identify the direction of systematic error in an investigation, and explain how to
reduce it using procedural issues in an experiment?
Can you show an example of every calculation that you use including headings, word
equations and numerical equations? This includes correct use of units and decimal places.
Can you gather and record data in a safe and accurate manner?
4
Lesson 2: The Mole, Mass and Avogadro’s Number Calculations [Top of Doc]
Learning Goals
UNDERSTANDING
- Can you define the term mole as a sample of particles, atoms, molecules, ions, etc. with 6.02 x 1023
species?
PROBLEM SOLVING
Can you determine the molar mass of a given compound?
Can you use molar mass and/or Avogadro’s number to convert mass, moles and number of
particles?
Can you use dimensional analysis to solve calculations involving masses, moles, number of
particles?
Dimensional Analysis Problems
Conversions Factors
1 hr = 60 min
1 min = 60 sec
1 ton = 2000 lbs
7 days = 1 week
24 hrs = 1 day
1 kg = 2.2 lbs
1 gal = 3.79 L
264.2 gal = 1 cubic meter
1 mi = 5,280 ft
1 kg = 1000 g
1 lb = 16 oz
20 drops = 1 mL
365 days = 1 yr
52 weeks = 1 yr
2.54 cm = 1 in
1 L = 1000 mL
0.621 mi = 1.00
km
1 yd = 36 inches
1 cc is 1 cm3
1 mL = 1 cm3
DIRECTIONS: Solve each problem using dimensional analysis. Every number must have a unit. Work
must be shown. Conversion factors are given below
1.) How many miles will a person run during a 10 kilometer race?
2.) The moon is 250,000 miles away. How many feet is it from earth?
3.) A family pool holds 10,000 gallons of water. How many cubic meters is this?
4.) The average American student is in class 330 minutes/day. How many hours/day is this?
How many seconds is this?
5
5) How many seconds are there in 1 year?
6) Lake Michigan holds 1.3 x 1015 gallons of water. How many liters is this?
7) Pepsi puts 355 ml of pop in a can. How many drops is this?
How many cubic meters is this?
8) Chicago uses 1.2 x 109 gallons of water /day. How many gallons per second must be pumped from the
lake every second to supply the city?
9) Sixty miles/ hour is how many ft/sec?
10) Lake Michigan holds 1.3 x 1015 gallons of water. If just Chicago removed water from the lake and it
never rained again, how many days would the water last? Chicago uses 1.2 x 109 gallons of water /day
11) How many minutes are in 180.0 days?
12) If a person weighs 125 lbs, 8 oz., how many mg does s/he weigh?
6
13) The distance from Santa Maria to Los Alamos is 16.25 mi. What is the distance in cm?
14) Santa Maria has an elevation of 6.30 x 105 mm. How many km is this elevation?
15) If a projectile travels 3.00 x 103 feet in one second, how far will it travel in 18 minutes?
16) A small herd of cattle consumes fourteen bales of hay in two weeks. How many bales will this herd
consume in a year?
17) During the previous year, Zach's weather station measured 0.8 yards of rain. Express this amount in cm.
18) If a swimmer swims 85.4 yards in five minutes, how many meters will s/he swim in 70.0 seconds?
19) Saffron costs $368.00 per ounce. Determine how many grams you can purchase for $15.00.
20) How many grams are equivalent to 1.80 x 10-4 tons? (English tons)
7
21). A gas station is charging $1.299 per gallon of gas. What would be the price for a liter of gas?
22). Determine the number of years in 8.35 x 106 minutes.
23). A quart of a liquid has a mass of 2.70 kilograms. How many quarts will take to weigh 100.0 pounds?
24). Sixty-two months is equivalent to how many seconds?
25). A car consumes 25.00 gallons of fuel when driving a distance of 400.0 km. How many gallons will it
consume when driving 250.0 miles?
26). 0.0054 weeks is equivalent to how many minutes?
27). How many feet per second is a wave going if it travels a distance of one mile in 7.35 seconds?
Molar Mass Calculations
8
9
Calculations continued on the following page.
10
Moles and Mass Calculations
Calculation of the number of moles of material in a given mass of that material
11
12
Moles, Mass and Avogadro’s Constant Calculations
1. What is the mass of 0.700 moles of Li2SO4 taking its molar mass to be exactly 110 g mol-1?
2. 0.200 moles of a substance has a mass of 27.0 g. What is the molar mass of the substance?
3. One drop of water weighs 0.040g. How many molecules are there in one drop, taking the
molar mass of water as exactly 18 g mol-1?
4. What is the mass (in g) of one molecule of sulphuric acid (H2SO4)?
5. A polymer molecule has a mass of 2.5 x 10-20g. What is the molar mass of the polymer?
7. How many moles are there in
6. Calculate (correct to 3 significant
figures) the mass of
a. 28.1g of silicon
a. 3.00moles of ammonia
b. ¼ mole of Li2O
b. 303g of KNO3
c. 0.0500 moles of aluminum nitrate
c. 4000 g of nickel sulphate
d. 3.01 x 1023 molecules of PCl3
d. 87.3g of methane?
e. 2.60 x 1022 molecules of
dinitrogen monoxide
13
8. 0.30 moles of a substance has a mass of 45g. What is its molar mass?
9. 3.01 x 1025 molecules of a gas have a mass of 6.40 kg. What is its molar mass?
10. Some types of Freon are used as a propellant in spray cans of paint, hair spray and other
consumer products. However, the use of Freons is being curtailed, because there is some
suspicion that they may cause environmental damage. If there are 25.00g of the Freon
CCl2F2 in a spray can, how many molecules are you releasing to the air when you empty the
can?
11. Vitamin C, ascorbic acid, has the formula C6H8O6.
a. The recommended daily dose of vitamin C is 60.0 milligrams. How many moles are
you consuming if you ingest 60 milligrams of the vitamin?
b. A typical tablet contains 1.00g of vitamin C. How many moles of vitamin C does this
represent?
c. When you consume 1.00g of vitamin C, how many oxygen atoms are you eating?
14
Extra 2 Step Conversion Problems
15
Lesson #3: Empirical and Molecular Formulas
[Top of Doc]
Learning Goals:
PROBLEM SOLVING
Can you determine empirical and molecular formulae using experimental data?
Empirical Formula
Find the empirical formula of the compound from the following data.
1. Ca 40%
C 12%
O 48%
Ca
C
O
Na
S
O
Mass
Moles
Ratio
Whole number ratio
Empirical Formula
2. Na 32.4%
S 22.5%
O 45.1%
Mass
Moles
Ratio
Whole number ratio
Empirical Formula
16
3. Na 29.1%
S 40.5%
O 30.4%
Na
S
O
Mass
Moles
Ratio
Whole number ratio
Empirical Formula
4. Pb 92.8%
O 7.20%
Pb
O
Pb
O
Mass
Moles
Ratio
Whole number ratio
Empirical Formula
5. Pb 90.66%
O 9.34%
Mass
Moles
Ratio
Whole number ratio
Empirical Formula
17
6. H 3.66%
P 37.8%
O 58.5%
H
P
O
Mass
Moles
Ratio
Whole number ratio
Empirical Formula
Empirical and Molecular Formula
Find the empirical and molecular formula of the compound using the following data.
Please layout your calculations in the way that you have been taught.
1.
C
85.7%
H
14.3%
Mr =
28 g mol-1
2.
C
85.7%
H
14.3%
Mr =
42 g mol-1
18
3.
P
10.88%
I
89.12%
4.
N
12.28%
H
3.51%
5.
P
43.66%
O
56.34%
6.
C
40%
H
6.67%
S
O
28.07%
53.3%
O
56.14%
Mr =
570 g mol-1
Mr =
228 g mol-1
Mr =
284 g mol-1
Mr =
60 g mol-1
19
Empirical and Molecular Formula from Experimental Data
Please make sure that you layout the calculations in the way that you have been taught.
1. 22.3g of an oxide of lead produced 20.7g of metallic lead on reduction with hydrogen
(Reduction is the removal of oxygen). Calculate the empirical formula of the oxide concerned.
PbxOy(s) + H2(g) → Pb(s) + H2O(l)
2. When 1.17g of potassium is heated in oxygen, 2.13g of an oxide is produced. What is the
empirical formula of the oxide produced? If the oxide has a molar mass of 71 g mol-1, what is
its molecular formula?
3. A hydrocarbon containing 92.3% of carbon has a molar mass of 26 g mol-1. What is the
molecular formula of the hydrocarbon?
20
4. When 1.335g of a chloride of aluminium is added to excess silver nitrate solution, 4.305g of
silver chloride is produced. Calculate the empirical formula of the chloride of aluminium. Hint;
you will need to work out how much chlorine there is in 4.305g of AgCl. This will be the
amount of chlorine in the initial 1.335g of the aluminium chloride.
5. 16g of a hydrocarbon burns in excess oxygen to produce 44g of carbon dioxide. What is the
empirical formula of the hydrocarbon? Hint; you will need to work out what mass of carbon is
contained in 44g of CO2. This is the mass of carbon in 16g of the hydrocarbon.
6. When the chloride of phosphorus containing 85.1% chlorine is heated, a second chloride
containing 77.5% chlorine is produced. Find the formula of the chlorides.
21
Lesson 4: Hydrates and Balancing Equations Review
[Top of Doc]
Learning Goals
PROBLEM SOLVING
- Can you determine the formula of a hydrate from experimental data?
UNDERSTANDING
- Do you know all of the polyatomic ions?
- Do you know the names and formulae of common compounds—water, ammonia,
hydrochloric acid, sulphuric acid, nitric acid, etc.?
- Do you know the oxidation states of the main elements?
- Do you know oxidation states for common transition elements—Ag+, Zn2+, Pb2+/Pb4+
EXPLAINING
Can you write names and formulae for ionic and covalent compounds?
Can you write symbolic equations from word equations?
Can you write word equations from symbolic equations?
Can you balance chemical equations?
Hydrates
In general physical changes do not create any new substances because no chemical bonds are
broken or formed. They are (in theory) reversible. Chemical changes always involve the breaking
and/or forming of new bonds, creating a product that is fundamentally different than its reactants.
Many ionic compounds naturally contain water as part of the crystal lattice structure.
A hydrate is a compound that has one or more water molecules bound to each formula unit.
The hydrated form of cobalt (II) chloride contains six water molecules in each formula unit as
shown above. The name of the compound is cobalt (II) chloride hexahydrate. The ‘hexa’ prefix is
used to show that there are 6 water molecules surrounding cobalt (II) chloride. Notice how cobalt is
a transition metal and the roman numeral (II) is used to show that it has an ionic charge of 2+.
When writing the formula of a hydrate, you can set water apart from the rest of the formula with a
dot. The coefficient is then placed after the dot to show how many water molecules are present.
A hydrate can usually be converted to the anhydrous form compound by heating which results in a
distinct colour change.
Learn more about hydrates: Hydrate Intro Video (mass and colour changes)
22
The Formula of a Hydrate
1. A 15.67 g sample of a hydrate of magnesium carbonate (MgCO3) was heated, without
decomposing the carbonate, to drive off the water. The mass was reduced to 7.58 g. What is
the formula of the hydrate?
2. A hydrate of Na2CO3 has a mass of 4.31 g before heating. After heating, the mass of the
anhydrous compound is found to be 3.22 g. Determine the formula of the hydrate and then
write out the name of the hydrate.
3. Given that the molar mass of Na2SO4 · nH2O is 322.1 g/mol, calculate the value of n.
4. A 5.00 g sample of hydrated barium chloride, BaCl2 · nH2O, is heated to drive off the water.
After heating, 4.26 g of anhydrous barium chloride, BaCl2, remains. What is the value of n in
the hydrate's formula?
23
Balancing Chemical Equations
1.
Ca
+
O2
CaO
2.
C
+
O2
CO2
3.
Fe
+
O2
Fe2O3
4.
H2O
H2
+
O2
5.
KCl
K
+
Cl2
+
Br2
6.
NaBr
+
Cl2
NaCl
7.
Mg
+
O2
MgO
BaCl2
Ba
+
Cl2
+
AgNO3
Cu(NO3)2
+
Ag
10. BeF2
+
Mg
MgF2
+
Be
11. KBr
+
Cl2
KCl
+
Br2
12. NaOH
+
FeCl2
NaCl
+
Fe(OH)2
8.
9.
Cu
24
13. AgNO3
+
Al2(SO4)3
Al(NO3)3
+
Ag 2SO4
14. K2SO4
+
Ca(NO3)2
KNO3
+
CaSO4
15. KCl
+
Pb(NO3)2
KNO3
+
PbCl2
16. Ca(OH)2 +
Fe(NO3)3
Fe(OH)3
+
17. MgSO4
+
Na3PO4
Mg3(PO4)2
+
Na2SO4
18. Na
+
H2O
NaOH
+
H2
19. Fe
+
H2O
Fe2O3
+
H2
20. CuCl2
+
H2S
CuS
+
HCl
Ca(NO3)2
Chemical Equations Practice
Write balanced equations from the following word equations. Make sure that you include
state symbols.
Sodium
Hydroxide
Solution
Copper (II)
Sulphate
Solution
Hydrochloric
+
Acid
+
Potassium
Hydroxide
Solution
→
→
Sodium
Chloride
Solution
Copper (II)
Hydroxide
Solid
+
+
Water
Potassium
Sulphate
Solution
25
Calcium
+ Water
Copper
Carbonate
Solid
+
Methane
Gas
+ Oxygen
Sulphuric
Acid
→
Calcium
Hydroxide
Solution
+
Hydrogen
→
Copper (II)
Sulphate
Solution
+
Water
→
Carbon
Dioxide
+
Water
+
Carbon
Dioxide
More Chemical Equations Practice
Write balanced equations from the following word equations. Make sure that you include
state symbols.
Potassium
Hydroxide
Solution
Lead (II)
Nitrate
Solution
Sulphuric
+
Acid
+
Potassium
Iodide
Solution
→
→
Potassium
Sulphate
Solution
Lead (II)
Iodide Solid
+
+
Water
Potassium
Nitrate
Solution
26
Calcium
Carbonate
Solid
Hydrochloric
+
Acid
Iron (III)
Oxide
Solid
Carbon
+ Monoxide
Gas
Ammonia
+ Nitric Acid
→
Calcium
Chloride
Solution
+
Water
→
Iron
+
Carbon
Dioxide
→
Ammonium
Nitrate
Solution
+
Carbon
Dioxide
Writing Chemical Equations from Experimental Data
1. In an experiment, a solution containing 3.31g of lead (II) nitrate reacts with a solution
containing 1.17g of sodium chloride to produce 2.78g of lead (II) chloride solid and leave a
solution that contains 1.70g of sodium nitrate. What is the equation for the reaction?
Equation:
………………………………………………………………………………………………..............................
27
2. In an experiment, a solution containing 6.675g of aluminium chloride reacted with a solution
containing 25.50g of silver nitrate. 21.52g of silver chloride was produced together with a
solution of 10.65g of aluminium nitrate. What is the equation for the reaction taking place?
Equation:
………………………………………………………………………………………………..............................
3. 1.133g of silver nitrate (silver is +1) was heated in an open tube. The silver residue weighed
0.720g. During the reaction 0.307g of nitrogen dioxide was also produced. The rest of the
mass loss was due to oxygen. Use the data to write the equation for the reaction.
Equation:
………………………………………………………………………………………………..............................
4. When 12.475g of hydrated copper (II) sulphate, CuSO4.xH2O was heated, 7.980g of
anhydrous salt was produced. Use the data to find the value of x and hence write the
equation for the reaction.
Equation:
………………………………………………………………………………………………..............................
5. When 13.9g of FeSO4.xH2O are heated, 4g of solid iron (III) oxide are produced together with
the loss of 1.6g of sulphur dioxide (SO2) and 2.0g of sulphur trioxide (SO3). The rest of the
mass loss being due to the water of crystallisation being lost. Use the data to write the full
equation for the action of heat.
Equation:
………………………………………………………………………………………………..............................
28
Lesson #5: Formative Lab 1 - Empirical Formulas
[Top of Doc]
Learning Goals:
PROBLEM-SOLVING
- Can you determine empirical and molecular formulas using experimental data?
EXPERIMENTAL SKILLS
- Can you create and/or use a correct data table? This includes correct labels, units and consistent use of
decimal places.
- Can you show an example of every calculation that you use including headings, word equations and
numerical equations? This includes correct use of units and decimal places.
- Can you identify the direction of systematic error in an investigation, and explain how to reduce it using
procedural issues in an experiment?
The Empirical Formula of Copper Sulfide
Objectives
● Make precise measurements of mass with an analytical balance.
● Determine the copper/sulfur mass ratio and the copper/sulfur mole ratio for a copper
sulfide produced by direct combination of the elements.
● Use these data to determine the simplest formula of the compound.
Introduction
Copper forms two sulfides, Cu2S and CuS. In this experiment you will make a sulfide of copper by
heating a known amount of copper with an excess amount of sulfur. During the heating, copper will
directly combine with sulfur to form a sulfide. By subtracting the mass of the copper used from the
mass of the compound formed, you can find the mass of the combined sulfur. From the masses of
copper and sulfur, the moles of each element, you can calculate the mole ratio and the empirical
formula.
Method
1. Measure the mass of the crucible and cover to the nearest mg (0.001 g).
2. Add a 10 cm piece of copper wire to the crucible. Determine the mass of the crucible, cover,
and copper.
3. Add enough sulfur to completely cover the copper wire. (It is not necessary to measure the
mass of crucible, cover, copper and sulfur since not all of this sulfur will react with the
copper.)
4. Heat the crucible, cover and contents with your burner in the hood. Heat them carefully for a
few seconds and then strongly until there is no further reaction. The excess sulfur, which
does not react with the copper, combines with oxygen in the air and departs the crucible as
sulfur dioxide.
5. Remove the crucible from the ring stand and allow the crucible to cool for 5 minutes
6. When the crucible has cooled, determine the mass of the crucible, cover, and contents.
Discard the contents of the crucible, clean the crucible.
Copper Sulfide Lab Video (7 minutes) (watch to record data)
Reaction of Copper and Sulfur (to observe video of the reaction in progress, watch from 1:30 to 3:00)
Quantitative Data
Unit and
Uncertainty
Mass of empty crucible and cover
Mass of crucible, cover, and
copper
29
Mass of crucible, cover, and
compound
Mass of copper used
Mass of compound formed
Mass of sulfur combined
Qualitative Data (Use before reaction, during reaction, and after reaction table)
Before Reaction
During Reaction
After Reaction
Processed Data
Cu
S
Mass
Moles
Copper/sulfur mass ratio
Copper/sulfur mole ratio
Empirical formula of copper
sulfide
Calculations (REMEMBER: Only 1 example of each type!)
30
Application of Principles
The actual formula of the copper sulphide formed should be CuS.
1. Suppose the mass of copper sulphide is larger than it should be at the end of your
experiment. What does that tell you about your ratio of Cu to S?
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
2. Suppose the mass of copper sulphide is smaller than it should be at the end of your
experiment. What does that tell you about your ratio of Cu to S?
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
3. Determine how each of the following errors would affect the mass of copper sulphide at the
end of the experiment and how each would affect the ratio of Cu to S.
a. The excess sulphur was not completely burned off at the end of the experiment.
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
b. Some of the copper did not react with the sulphur.
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
c. Some of the sample spatters out of the crucible when the excess sulphur is being
burned off.
31
……………………………………………………………………………………………………………..
…………………………………………………………………………………………………………….
d. The copper reacts with oxygen in the air instead of sulphur.
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
4. Which of the errors listed in question 3 were most likely to occur in your trial? Explain.
……………………………………………………………………………………………………………..
……………………………………………………………………………………………………………..
Practice
1. If 52.2 g of bismuth react under certain conditions with 10.0 g of oxygen, what is the empirical
formula of the bismuth oxide formed?
Bi
O
Mass
Moles
Bismuth/oxygen mole ratio
Empirical formula of bismuth
oxide
2. 59.9 g of titanium metal react completely with 40.1 g of sulfur. What is the empirical formula of
the titanium sulfide.
Ti
S
Mass
Moles
Titanium/sulfur mole ratio
Empirical formula of titanium
sulfide
32
Lesson #6: Mass-Mass and Mass-# Relationships
[Top of Doc]]
Learning Goals
UNDERSTANDING
- Do you know the difference between theoretical and actual yield?
EXPLAINING
- Can you write symbolic equations from word equations?
- Can you balance equations?
PROBLEM SOLVING
- Can you determine the molar mass of a given compound?
- Can you solve problems involving chemical equations and mass?
- Can you create ratios between the coefficients in balanced equations to solve unknown values?
- Can you calculate percent yield based on actual yield and calculated theoretical yield?
Calculation of Products/Reactants Based on Equations
1. What mass of barium sulphate would be produced from 10g of barium chloride in the following
reaction?
BaCl2 + H2SO4 → BaSO4 + 2HCl
2. What mass of potassium chloride would be produced from 20g of potassium carbonate in the
following reaction?
K2CO3 + 2HCl → 2KCl + CO2 + H2O
3. A solution of copper sulphate reacts with sodium hydroxide solution to produce a precipitate of
copper hydroxide according to the following equation
CuSO4(aq) + NaOH(aq) → Cu(OH)2(s) + Na2SO4(aq)
What mass of sodium hydroxide would be needed to convert 15.95g of copper sulphate to
copper hydroxide and what mass of copper hydroxide would be produced?
33
4. In a reaction between calcium carbonate and nitric acid, what mass of calcium nitrate would
be produced from 33.3g of calcium carbonate?
5. In the following reactions, calculate the mass of precipitate formed from 20g of zinc (II)
sulfate in each case.
a. ZnSO4(aq)
+
2NaOH
→ Zn(OH)2(s)
+
Na2SO4(aq)
ai) Calculate the number of aqueous cations produced from 20g of aluminum sulfate.
b. Al2(SO4)3(aq) +
6NaOH
→ 2Al(OH)3s)
+
3Na2SO4(aq)
bi) Calculate the number of aqueous cations produced from 20g of magnesium sulfate.
c. MgSO4(aq)
+
2NaOH
→ Mg(OH)2(s)
+
Na2SO4(aq)
ci) Calculate the number of sulfate ions produced from 20g of magnesium sulfate.
34
Lesson #7: Limiting Reagents, Excess, Experimental and Percent
Yield [Top of Doc]
Learning Goals
UNDERSTANDING
Do you know the difference between theoretical and experimental yield?
EXPLAINING
Can you balance equations?
Can you write symbolic equations from word equations?
Can you write word equations from symbolic equations?
PROBLEM SOLVING
Can you solve problems involving chemical equations and mass?
Can you calculate percent yield based on experimental yield and calculated theoretical yield?
Can you identify the limiting and excess reagent in a problem?
Can you solve problems involving limiting reagents?
Can you calculate the moles, mass or number of species in excess based on the amount of limiting reagent
used?
Limiting Reagents
For the following reactions, find the following:
a)
Which of the reagents is the limiting reagent?
b)
What is the maximum amount of each product that can be formed?
c)
How much of the other reagent is left over after the reaction is complete?
1)
Consider the following reaction:
3 NH4NO3 + Na3PO4 → (NH4)3PO4 + 3 NaNO3
Answer the questions above, assuming we started with 30 grams of ammonium nitrate and 50
grams of sodium phosphate.
2)
Consider the following reaction:
3 CaCO3 + 2 FePO4 → Ca3(PO4)2 + Fe2(CO3)3
Answer the questions at the top of this sheet, assuming we start with 100 grams of calcium
carbonate and 45 grams of iron (III) phosphate.
35
More Problems on Theoretical Yield and Limiting Reagent
1. The reaction of hydrogen with oxygen to produce water is described by the following equation:
2 H2 (g) + O2 (g) → 2 H2O (l)
You react 150 H2 molecules with 100 O2 molecules to produce H2O.
a. Which is the limiting reactant, hydrogen or oxygen?
b. How many water molecules can you produce from your supply of hydrogen and oxygen?
2. If 6 moles of hydrogen (H2) and 4 moles of oxygen (O2) are mixed and reacted, which is the
limiting reactant? How many moles of water would be produced?
3. If you had 1.73 moles of hydrogen (H2) and 0.89 moles of oxygen (O2) which is the limiting
reactant? How many moles of water can you produce from your supply of hydrogen and
oxygen?
36
4. If you had 17.3 g of hydrogen gas and 8.91 g of oxygen gas, which is the limiting reagent and
how many grams of water could you produce?
5. Hydrogen cyanide is used in the production of cyanamid fertilizers. It is produced by the
following reaction:
2 CH4 (g) + 2 NH3 (g) + 3 O2 (g) → 2 HCN (g) + 6 H2O (l)
How much hydrogen cyanide can be produced starting with 100 kg of each of the reactants?
6. When 5.00 g of KClO3 is heated it decomposes according to the equation:
2 KClO3 (s) → 2 KCl (s) + 3 O2 (g)
a)
Calculate the theoretical yield of oxygen.
b)
Give the percent yield if 1.78 g of oxygen is produced.
c)
How much oxygen would be produced if the percentage yield was 78.6%?
7. What is the percent yield of water if 138 g of water is produced from 16 g of hydrogen and an
excess of oxygen? [Hint: Write the equation.]
37
8. What is the percent yield of ammonia (NH3) if 40.5 g of ammonia is produced from 20.0 mol of
hydrogen and an excess of nitrogen? [Equation?]
9. What is the percent yield of water if 58 g of water are produced by combining 60 g of oxygen
and 7.0 g of hydrogen?
10. The electrolysis of water from H2 and O2:
2 H2O (l) → 2 H2 (g) + O2 (g)
What is the percent yield of oxygen if 12.3 g of oxygen is produced from the decomposition of
14.0 g of water?
11. 107 g of oxygen is produced by heating 300 g of potassium chlorate (KClO3). Calculate the
percent yield.
38
12. What is the percent yield of iron (II) sulfide if 3.00 moles of iron react with excess sulfur to
produce 220 g of iron (II) sulfide?
Fe (s) + S (s) → FeS (s)
13. Iron pyrites (FeS2) react with oxygen according to the following equation:
4 FeS2 (s) + 11 O2 (g) → 2 Fe2O3 (s) + 8 SO2 (g)
If 300 g of iron pyrite is burned in 200 g of oxygen, 143 grams of iron (III) oxide is produced.
What is the percent yield of iron (III) oxide?
14. 70 g of manganese (IV) oxide is mixed with 3.5 moles of hydrochloric acid. How many grams
of chlorine gas (Cl2) will be produced from this reaction if the percent yield for the process is
42%?
39
Limiting Reagents and Percentage Yield
1. Consider the reaction:
H2O2(aq) + 2KI(aq) + H2SO4(aq) → I2(s) + K2SO4(aq) + 2H2O(l)
What mass of iodine is produced when 100.00g of KI is added to a solution containing 12.00g
of H2O2 and 50.00g of H2SO4?
2. Consider the reaction:
2Al(s) + 3I2(s) → 2AlI3(s)
Determine the limiting reagent and the theoretical yield from:
a. 1.20mol Al and 2.40 mol iodine
b. 1.20g Al and 2.40g iodine
c. How many grams of Al are left over in part b.
40
3. Freon-12 (used as a coolant in refrigerators), is formed as follows:
3CCl4(l) + 2SbF3(s) → 3CCl2F2(g) + 2SbCl3(s)
150g CCl4 is combined with 100g SbF3 to give Freon-12 (CCl2F2).
a. Identify the limiting and the excess reagents.
b. How many grams of Freon-12 can be formed?
c. How much of the excess reagent is left over?
in grams?
number of molecules?
4. Aspirin is made by adding acetic anhydride to an aqueous solution of salicylic acid. The
equation for the reaction is:
2C7H6O3(aq) + C4H6O3(l) → 2C9H8O4(aq) + H2O(l)
If 1.00kg of salicylic acid is used with 2.00kg of acetic anhydride, determine:
a. the limiting reagent
b. the theoretical yield of aspirin
41
c. If 1.12kg aspirin is produced experimentally, what is the percentage yield?
Limiting Reagents and Percentage Yield #2
[Extra Practice]
1.
2.
3.
4.
42
5.
6.
7.
43
Lesson #8: Formative Lab 2 - Limiting Reagents
[Top of Doc]
Learning Goals
PROBLEM SOLVING
Can you solve problems involving chemical equations and mass?
Can you calculate percent yield based on experimental yield and calculated theoretical
yield?
Can you identify the limiting and excess reagent in a problem?
Can you solve problems involving limiting reagents?
EXPERIMENTAL SKILLS
Can you create and/or use a correct data table? This includes correct labels, units and
consistent use of decimal places.
Can you show an example of every calculation that you use including headings, word
equations and numerical equations? This includes correct use of units and decimal places.
Can you gather and record data in a safe and accurate manner?
Can you identify the direction of systematic error in an investigation, and explain how to
reduce it using procedural issues in an experiment?
Limiting Reagents Formative Lab
Objective: To evaluate your ability to apply your knowledge of the following content areas:,
balancing chemical equations, stoichiometry calculations (mole conversions), and limiting reactant
analysis.
Part A: Getting Started
Today’s experiment is based on this reaction of aluminium with copper (II) chloride, which will give
copper and aluminium chloride.
1. Write a balanced equation for the reaction.
……………………………………………………………………………………………………………
Part B: Pre-Lab Calculations
In today’s lab you will be reacting a 0.15 g piece of aluminum wire with 20.0 mL of copper (II) chloride
solution (containing 0.53 grams CuCl2 in water). Use these two given reactant quantities, 0.15 grams
of Al and 0.53 grams of CuCl2 to determine the mass of Cu(s) product that should be formed (the
theoretical yield). This is a limiting reactant problem using two givens.
2. What mass of copper, Cu, should theoretically be formed? (theoretical yield)
44
3. Which of the two reactants was the limiting reactant? __________________
4. At the end of the reaction, how much of the excess reactant will be left over in grams?
Part C: Experimental Procedure
Caution you MUST wear your safety goggles during the rest of the lab. Do not take them off until you
have cleaned up your lab table and returned to your seat!
Step 1 Obtain a preweighed piece of aluminium wire and measure out 20.0mL of CuCl2(aq)
solution and pour it into a clean, dry 50mL beaker. The color should be blue because of the
Cu+2 ions floating in the solution.
Step 3. Coil the aluminium wire and place it into the beaker so that it is immersed into the
solution.
Step 4. Allow the reaction to occur for 10 minutes and record your detailed observations of
everything BEFORE, DURING, and AFTER the reaction in table A below. Include
observations of size, colour, texture, temperature, bubbles, etc., but not taste!
Step 5: Weigh a piece of filter paper. Place into a filter funnel.
Step 6. After the reaction is complete, pour the aqueous mixture into the funnel with filter paper.
This will leave the solid component in the beaker. Remove the solid copper part from the
excess aluminum wire by using a glass rod and shots of distilled water. Decant the copper
components onto the filter paper. Rinse the beaker with distilled water a few times to get all
remaining copper onto the filter paper. NOTE: The excess aluminum wire should not be
filtered (it stays in the reaction beaker).
Step 7. Allow the filter paper to dry in the oven at the back of the room and then reweigh. Place
your waster in the correct container (as advised by your teacher).
Limiting Reagent Lab (8 minutes) (Extra Support)
Limiting Reagents Answer Checker
Table #__________: Add Descriptive Title…………
BEFORE
Appearance of Al and CuCl2
DURING THE RXN
AFTER
45
,
Mass of CuCl2(aq):
Mass of Al (s) :
Initial Mass of Filter Paper : ………………………………….
Final Mass of Filter Paper + Cu: : ………………………………….
Final mass of copper: ………………………………….
Percent Yield: ………………………………………….. (Direction of Error - above or below expected?)
5. Was your final mass larger or smaller than expected? Use your observations to explain why
this might be the case.
……………………………………………………………………………………………………………
……………………………………………………………………………………………………………
……………………………………………………………………………………………………………
……………………………………………………………………………………………………………
……………………………………………………………………………………………………………
……………………………………………………………………………………………………………
……………………………………………………………………………………………………………
6. What improvements could you make to this laboratory based on your answer to the previous
question?
……………………………………………………………………………………………………………
46
……………………………………………………………………………………………………………
……………………………………………………………………………………………………………
……………………………………………………………………………………………………………
……………………………………………………………………………………………………………
Weakness or Error
Explanation of Effect of
Weakness or Error on Your
Results
Valid Improvement or
Solution to Identified
Weakness or Error
Lesson #10: Unit 1 - Moles and Masses Review
[Top of Doc]
Moles and Masses Review Questions
I recommend that you do one from each of the sections rather than doing each section in order.
Moles to particles
1.
2.
3.
4.
5.
How many As atoms are there in 3.68 moles of As?
How many moles are there in 6.74×1024 atoms?
How many molecules of nitrogen dioxide are there in 4.02 moles?
How many oxygen atoms will be present in the molecules in question three?
How many nitrate ions are there in 0.713 moles of calcium nitrate?
Molar masses
Find the molar mass of the following and identify each as being either molecular, atomic or formula
masses.
1.
2.
3.
4.
5.
Au
NH4Cl.4H2O
C6H12O6
H2
Pb
Moles-Mass conversion
1.
2.
3.
4.
What is the mass of 6.07 moles of sodium fluoride?
How many moles are represented by 4.96 g of sodium dichromate?
What mass of octane, (C8H18) does it take to have 5.00 moles?
How many moles are in 43.1 g of potassium chromate dihydrate? How many moles of water is
that?
47
Limiting reagent problems
1. 30.0 g of Carbon reacts with 15.0 g of oxygen to form Carbon monoxide in an incomplete
combustion
a. What is the limiting reagent?
b. What is the theoretical yield of product?
c. How many grams of excess reagent remain?
2. 320 g of sulphur dioxide react with 32.0 g of oxygen and excess water to form sulphuric acid,
H2SO4
a. What is the limiting reagent?
b. What is the theoretical yield of product?
c. How many grams of excess reagent remain?
Empirical formula
1. If 3.50 g of iron reacts with 1.50 g of oxygen, what is the empirical formula of the compound
created?
2. Analysis shows a compound to contain 25.56% potassium, 35.41% chromium and 36.73%
oxygen. What is it empirical formula?
3. Analysis of 20.0 g of a compound known to contain only calcium and bromine indicates that
4.00 g of calcium are present. Find the empirical formula.
Molecular formula
1. A compound contains 56.3% oxygen and 43.7% phosphorus. The molar mass of the
compound was determined by mass spectroscopy to be 283.89 g mol-1. What is its molecular
formula?
2. Determine the molecular formula of a compound with empirical formula CH and a molar mass
of 78.11 g mol-1.
3. If 4.04 g of nitrogen combine with 11.54 g of oxygen to produce a compound with a molar
mass of 108.0 g mol-1, what is the molecular formula?
Balancing equations
Balance the following equations:
1. C4H10(g) + O2(g) → CO2(g) + H2O(g)
2. Al(s) + Br2(l) → Al2Br6(s)
3. SiO2(s) + C(s) → Si(s) + CO(g)
4. UO2(s) + HF(g) → UF4(s) + H2O(l)
5. H3BO3(s) → B2O3(s) + H2O(l)
Stoichiometry
1. Magnesium reacts with nitrogen to form magnesium nitride
a. How many grams of magnesium nitride are formed from 11.2 g of magnesium?
48
b. You end up with 5.67 g of magnesium nitride. How many grams of nitrogen did you
have initially?
2. Potassium chloride reacts with oxygen to form potassium chlorate
a. How many moles of potassium chlorate are formed from 9.34 g of potassium chloride?
b. If 34.6 g of potassium chlorate are produced, how many moles of oxygen with there
initially?
Moles and Mass Example Test
Avogadro’s Number = 6 x 1023
1. The formula of ethyl ethanoate is CH3CO2C2H5.
a. How many molecules are in 0.65 moles?
(1)
b. How many atoms are present in 0.65 moles of molecules?
(2)
c. How many carbon atoms are present in 0.65 moles?
(2)
2. (NH4)3PO4 is an ionic compound.
a. List the formula of the ions present in (NH4)3PO4?
(2)
b. How many positive ions are present in 0.76 moles?
(2)
c. How many hydrogen atoms are present in 0.76 moles?
(2)
3. When 5.00 g of FeCl3 · xH2O are heated, 2.00 g of H2O are driven off. Find the value of x
(3)
49
4. Strychnine, a deadly poison, has a molar mass of 334 g mol-1. 100 g of strychnine contains
75.42 g carbon, 6.63 g hydrogen, 8.38 g nitrogen, and the balance oxygen. What is the
molecular formula of strychnine?
(4)
5. What mass of K2Cr2O7 contains 0.035 moles?
6. Find mass of 3.16 x 10 12 molecules of AlCl3.
(1)
(2)
7. Calcium carbonate reacts with nitric acid (HNO3) to produce calcium nitrate, carbon dioxide
and water.
a. Write a balanced equation for the reaction.
(2)
b. Find theoretical yield of calcium nitrate if 4.5 g of nitric acid was added to an excess of
calcium carbonate.
(3)
c. The mass of calcium nitrate produced was 0.89 g. Find the % yield.
(1)
50
8. Iron (II) sulphate reacts with sodium hydroxide to form sodium sulphate and iron (II) hydroxide.
a. Write a balanced equation for the reaction.
(3)
b. 15 g of iron (II) sulphate are added to 7 g of sodium hydroxide. Find the limiting
reactant.
(3)
c. Find the mass of sodium sulphate produced in the reaction.
(2)
d. Find the mass remaining of the excess reactant.
(3)
51
Unit 2: Gases and Solutions - Learning Goal Summary
Can You? / Do You?
UNDERSTANDING
Do you understand that equal volumes of gases, under the same conditions of temperature and
pressure, contain the same number of particles? (Avogadro’s Law)
Do you understand the four assumptions of kinetic molecular theory (KMT)?
EXPLAINING
Can you explain gas phenomena (pressure, temperature, volume conditions or changes) using the ideas
of kinetic molecular theory?
PROBLEM SOLVING
Can you solve problems related to Avogadro's Law (using volumes of gases and stoichiometric analysis,
including the concept of limiting reagents)?
Can you Convert Celsius and Kelvin temperatures?
Can you use the combined gas law to solve problems concerning pressure, volume and temperature?
Can you use the ideal gas law to solve problems concerning pressure, volume, temperature and moles
of gas?
Can you convert between pressure units? (1.0 atm = 101.3 kPa = 1.013 x 105 Pa)
Can you use the molar gas volume, taken as 22.7 dm3 at room temperature and pressure.
Do you understand that you must use kPa (pressure) and dm3 or L (volume) when using R = 8.314 J/mol
K?
Can you Uuse the ideal gas equation to calculate the molar mass of a gas, given the necessary
information?
EXPERIMENTAL SKILLS
Can you create and/or use a correct data table? This includes correct labels, units and consistent use of
decimal places.
Can you gather and record data in a safe and accurate manner?
Can you analyse your data and communicate it in an appropriate manner?
Uses critical thinking skills to evaluate information and drawing appropriate conclusions?
Can you evaluate weaknesses and errors in your method based on experimental evidence?
Lesson 12 - Unit 2 Lesson 1
Introduction to Ideal Gases and Kinetic Molecular Theory
[Top of Doc]
Learning Goals:
EXPLAINING
Can you explain gas phenomena (pressure, temperature, volume conditions or changes) using the ideas
of kinetic molecular theory?
UNDERSTANDING
Do you understand that equal volumes of gases, under the same conditions of temperature and
pressure, contain the same number of particles? (Avogadro’s Law)
Do you understand the four assumptions of kinetic molecular theory (KMT)?
PROBLEM SOLVING
Can you solve problems related to Avogadro's Law (using volumes of gases and stoichiometric analysis,
including the concept of limiting reagents)?
Avogadro’s Law: Gas Properties Simulation Demonstration
Introduction:
This simulation allows you to investigate the meaning and implications of Avogadro’s Law of gases:
52
Under the same conditions of pressure, volume and temperature
different samples of gas will have the same number of particles.
The simulation uses a screen that allows the
users to manipulate all of the variables that
affect gas properties.
As the variables get manipulated, your job is to
observe carefully, document your observations,
draw conclusions and ask questions if you have
them.
Part A: Effect of Mass of Gas (Gas Identity)
1. 50 “heavy particles” at T=300K and full volume: Allow the pressure to stabilize. What is the
pressure?
2. 50 “light particles” at T=300K and full volume: Allow the pressure to stabilize. What is the
pressure?
3. You can assume that the “heavy” particles were a gas like Xe. What is the molar mass of Xe?
4. You can assume that the “light” particle were a gas like He. What is the molar mass of He?
5. What effect does the size/identity of a gas have on the pressure (when number of particles,
temperature and volume do not change)?
Part B: Effect of Number of Particles (Moles of Gas)
6. Fill out the table below as your teacher changes the number of light gas particles.
Number of light gas particles
Pressure (Atm)
50
100
200
400
53
7. What effect does the number of gas particles have on the pressure (when volume and
temperature do not change)? Why?
Part C: Statement of Avogadro’s Law
Now play around with the volume and the temperature of the gas particles. Watch what happens to
the pressure during these manipulations.
8. Why does Avogadro’s Law have to specify “under the same conditions of pressure, volume
and temperature” in the statement of the law?
Kinetic Molecular Theory (KMT)
Avogadro's Law:
Under the same conditions of pressure, volume and temperature different samples of gas will have
the same number of particles. In other words…
ALL GASSES BEHAVE EXACTLY THE SAME WAY...IDENTITY DOES NOT MATTER.
Kinetic Molecular Theory (KMT):
Postulate #1 is at the heart of Avogadro’s Law:
Explain why postulate number 1 means that 1 mol of hydrogen (H2; molar mass = 2 g/mol) and 1 mol
of butane (C4H10; molar mass = 58 g/mol) will have exactly the same pressure and volume at 298K
even though a molecule of butane is 29 TIMES bigger than a molecule of hydrogen.
54
Avogadro’s Law Problems
1. According to the equation below, what volume of nitrogen dioxide would you expect to be
formed from 20 cm3 of nitrogen monoxide, assuming that the volumes are measured at the
same temperature and pressure?
2NO(g) + O2(g) → 2NO2(g)
2. Four identical flasks are filled with hydrogen, oxygen, carbon dioxide and chlorine at the same
temperature and pressure. Which flask will have the greatest mass? Why?
3. A mixture of 20cm3 hydrogen and 40cm3 oxygen is exploded in a strong container. After
cooling to the original temperature and pressure (at which water is a liquid) what gas, if any,
will remain in the container?
4. What volume of sulphur trioxide would be produced by the complete reaction of 100cm3 of
sulphur dioxide with oxygen? What volume of oxygen would be needed to just react with
sulphur dioxide?
SO2(g)) + O2(g) → 2SO3(g)
5. In a reaction between methane and oxygen, 60 cm3 of methane was burnt with 60 cm3 of
oxygen. What is the composition of the gas mixture produced at room temperature?
55
Lesson 13 - Unit 2 Lesson 2
[Top of Doc]
The Combined Gas Law
Learning Goals:
PROBLEM-SOLVING
Can you use the combined gas law to solve problems concerning pressure, volume and temperature?
EXPERIMENTAL SKILLS
Can you analyze your data and communicate it in an appropriate manner?
Uses critical thinking skills to evaluate information and drawing appropriate conclusions?
PROBLEM SOLVING
Can you use the combined gas law to solve problems concerning pressure, volume and temperature?
EXPERIMENTAL SKILLS
Can you analyze your data and communicate it in an appropriate manner?
Can you use critical thinking skills to evaluate information and drawing appropriate conclusions?
Factors Affecting Gases Activities
This combination of teacher demonstration, student lab and online activities will allow you to
investigate the effect that changing volume, temperature or pressure has on the properties of a gas.
In every case one of the three variables (P, V or T) will be held constant while another is
manipulated. You will then observe the remaining variable to see how it is affected.
Student Investigation 1: Effect of Temperature on Volume of a Gas (Constant Pressure)
Watch the following “Coke can” demonstration video or your teacher’s demonstration in class.
1. Briefly describe what happened during the demonstration.
2. What are the conditions of temperature, volume and pressure at the beginning and end of the
demonstration? You can look up current pressure at
https://www.wunderground.com/id/jakarta. You have to divide hPa by 10 to get kPa.
56
Initial
Final
Temperature
High
Low
High
Low
Volume
High
Low
High
Low
Pressure (kPa)
3. What effect does changing temperature have on a fixed volume of gas at constant pressure?
Why?
Student Investigation 2 : Effect of Volume on Pressure (at constant temperature)
During this investigation you will not only see the effect that volume has on pressure, with constant
temperature, but you will also learn how to use LoggerPro and a Vernier Gas Pressure Sensor to
collect data.
Setting up the Vernier Probe and LoggerPro:
1. Open up LoggerPro, set up the LabQuest Mini and Gas Pressure Sensor with syringe as in
the following photograph.
2. Insert the USB plug into your computer. You will know that it has been successful when you
see a pressure readout in the lower left-hand corner of your screen.
57
3. Take the syringe off of the pressure sensor, put the plunger on 30 cm3, then re-attach it to the
Gas Pressure Sensor.
4. Go to Experiment, choose Data Collection and use the drop-down menu to choose Events
with Entry. Put the following labels and units in:
5. Click the green arrow to start data collection. When you are ready to take your first reading,
click on the circular camera lens icon (far right) and enter the volume (30) and click OK.
6. Reduce the volume by 2 cm3 and repeat the process. Keep doing this until it becomes
impossible to reduce the volume any more then click the red stop button.
7. Click on the autoscale button (the capital “A”) above the graph to scale your data.
8. Right-click inside the graph and use Graph Options to untick the Connect Points box.
9. Go to Analyze then Curve Fit…
10. Choose Linear and Try Fit. Look at the fit. Then choose Inverse (A/x) and Try Fit. Look at the
fit. Which is better? Click on OK then show your curve to your teacher.
11. After your teacher approves your curve fit, address the following questions:
What effect does reducing the volume of a sample of gas (at constant temperature) have on the
pressure of the gas? Why?
Why must the curve fit not be linear?
58
4. Read this presentation on using Logger Pro to become familiarized with making graphs on Logger
Pro.
Student Activity—Gas Laws Simulations:
Go to the “Factors Affecting Gases and Student Activity-Gas Laws Simulation” (also linked on the website)
and open the gas law simulation that you will find there)
Effect of Temperature on Volume (at constant pressure):
Use the Cool button to change the temperature of the gas sample. For each manipulation, record the
temperature and volume in the table below:
Temperature (K)
Volume (L)
Sketch a graph curve on the graph axes below that represents your data (it’s just “sketch”!). Explain the
line/curve.
How well does this fit with the “Coke can” demonstration earlier in this unit?
Explanation of trend:
59
Effect of Temperature on Pressure (at constant volume):
Reload the simulation by clicking the refresh button in your browser. Click the lock icon in the Volume box to
maintain a constant volume. Use the Heat button to increase the temperature of the gas sample. For each
manipulation, record the temperature and pressure in the table below:
Temperature (K)
Pressure (atm)
Sketch a graph curve on the graph axes below that represents your data (it’s just “sketch”!). Explain the
line/curve.
Explanation of trend:
Effect of Number of Gas Particles on Pressure (at constant volume):
Reload the simulation by clicking the refresh button in your browser. Click the lock icon in the Volume box to
maintain a constant volume. Click on the air pump handle to increase the number of gas particles in the
cylinder. Each pump puts 10 additional gas particles into the cylinder. For each manipulation, record the
number of particles and pressure in the table below:
Number of Particles
Pressure (atm)
60
Sketch a graph curve on the graph axes below that represents your data (it’s just “sketch”!). Explain the
line/curve.
Explanation of trend:
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The Gas Laws and Moles and Gases (Practice Problems)
Boyle’s Law
1)
If I have 5.6 dm3 of gas in a piston at a pressure of 1.5 atm and compress the gas until its volume is 4.8
dm3, what will the new pressure inside the piston be?
2)
I have added 15 dm3 of air to a balloon at sea level (1.0 atm). If I take the balloon with me to Denver,
where the air pressure is 0.85 atm, what will the new volume of the balloon be?
3)
I’ve got a car with an internal volume of 12,000 dm3. If I drive my car into the river and it implodes,
what will be the volume of the gas when the pressure goes from 1.0 atm to 1.4 atm?
Charles’s Law
1)
If I have 45 dm3 of helium in a balloon at 250 C and increase the temperature of the balloon to 550 C,
what will the new volume of the balloon be?
2)
Calcium carbonate decomposes at 12000 C to form carbon dioxide and calcium oxide. If 25 dm3 of
carbon dioxide are collected at 12000 C, what will the volume of this gas be after it cools to 250 C?
3)
I have 130 dm3 of gas in a piston at a temperature of 2500 C. If I cool the gas until the volume
decreases to 85 dm3, what will the temperature of the gas be?
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Combined Gas Law Worksheet
Consider watching Be Lazy Don’t Memorize the Gas Laws (7 minutes)
1)
If I initially have 4.0 dm3 of a gas at a pressure of 1.1 atm, what will the volume be if I increase the
pressure to 3.4 atm?
2)
A toy balloon has an internal pressure of 1.05 atm and a volume of 5.0 L. If the temperature where the
balloon is released is 200 C, what will happen to the volume when the balloon rises to an altitude
where the pressure is 0.65 atm and the temperature is –150 C?
3)
A small research submarine with a volume of 1.2 x 105 dm3 has an internal pressure of 1.0 atm and an
internal temperature of 150 C. If the submarine descends to a depth where the pressure is 150 atm
and the temperature is 30 C, what will the volume of the gas inside be if the hull of the submarine
breaks?
4)
People who are angry sometimes say that they feel as if they’ll explode. If a calm person with a lung
capacity of 3.5 dm3 and a body temperature of 360 C gets angry, what will the volume of the person’s
lungs be if their temperature rises to 390 C. Based on this, do you think it’s likely they will explode?
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Lesson 14 - Unit 2 Lesson 3
[Top of Doc]
The Ideal Gas Law, PV = nRT
Learning Goals:
PROBLEM-SOLVING
Can you convert Celsius and Kelvin temperatures?
Can you use the combined gas law to solve problems concerning pressure, volume and temperature?
Can you use the ideal gas law to solve problems concerning pressure, volume, temperature and moles
of gas?
Can you convert between pressure units? (1.0 atm = 101.3 kPa = 1.013 x 105 Pa)
Do you understand that you must use kPa (pressure) and dm3 or L (volume) when using R = 8.314 J/mol
K?
Can you use the ideal gas equation to calculate the molar mass of a gas, given the necessary
information?
Combined Gas Law Practice Sheet
1)
A bag of potato chips is packaged at sea level (1.00 atm) and has a volume of 315 cm3. If this bag of
chips is transported to Denver (0.775 atm), what will the new volume of the bag be?
2)
A Los Angeles class nuclear submarine has an internal volume of eleven million dm3 at a pressure of
1.250 atm. If a crewman were to open one of the hatches to the outside ocean while it was
underwater (pressure = 15.75 atm), what would be the new volume of the air inside the submarine?
3)
A child has a toy balloon with a volume of 1.80 dm3. The temperature of the balloon when it was filled
was 200 C and the pressure was 1.00 atm. If the child were to let go of the balloon and it rose 3
kilometers into the sky where the pressure is 0.667 atm and the temperature is -100 C, what would the
new volume of the balloon be?
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4)
A commercial airliner has an internal pressure of 1.00 atm and temperature of 250 C at takeoff. If the
temperature of the airliner drops to 170 C during the flight, what is the new cabin pressure?
5)
If divers rise too quickly from a deep dive, they get a condition called “the bends” which is caused by
the expansion of very small nitrogen bubbles in the blood due to decreased pressure. If the initial
volume of the bubbles in a diver’s blood is 15 cm3 and the initial pressure is 12.75 atm, what is the
volume of the bubbles when the diver has surfaced to 1.00 atm pressure?
Ideal Gas Law Practice Worksheet
Solve the following problems using the ideal gas law (Use R = 8.31 J mol-1 K-1 or 0.0821 L atm mol-1 K-1)
1)
How many moles of gas does it take to occupy 120 liters at a pressure of 233 kPa and a temperature of
340 K?
2)
If I have a 50 dm3 container that holds 45 moles of gas at a temperature of 2000 C, what is the pressure
inside the container?
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3)
It is not safe to put aerosol canisters in a campfire, because the pressure inside the canisters gets very
high and they can explode. If I have a 1.0 dm3 canister that holds 2 moles of gas, and the campfire
temperature is 14000 C, what is the pressure inside the canister?
4)
How many moles of gas are in a 30 dm3 scuba canister if the temperature of the canister is 573 K and
the pressure is 20,265 kPa?
5)
I have a balloon that can hold 100 dm3 of air. If I blow up this balloon with 3 moles of oxygen gas at a
pressure of 101 kPa, what is the temperature of the balloon?
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Lesson 15 - Unit 2 Lesson 4
[Top of Doc]
Deviations from Ideal Gas Behaviour
PROBLEM-SOLVING
Can you solve problems using molar gas volume, taken as 22.7 dm3/22.4 L at room temperature and
pressure?
Can you identify and explain the conditions for the greatest and lowest deviations from ideal gas
behaviours?
Can you analyze and interpret graphs to solve problems related to deviations from ideal gas behaviours?
Therefore, based on the assumptions above, conditions that 1) make intermolecular forces more
significant and 2) make the volume of each particle more significant will result in greater deviations in
ideal gas behaviour.
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Conclusion: Greatest Deviations at Low Temperature, Low Volume Environment.
Source: POGIL Activities for Chemistry
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Lesson 16 - Unit 2 Lesson 5
Formative Quiz 1/Gas Stoichiometry and Limiting Reagents
[Top of Doc]
Learning Goals:
PROBLEM-SOLVING
Convert Celsius and Kelvin temperatures?
Use the combined gas law to solve problems concerning pressure, volume and temperature?
Use the molar gas volume, taken as 22.7 dm3 at room temperature and pressure.
Use the ideal gas law to solve problems concerning pressure, volume, temperature and moles of gas?
Convert between pressure units? (1.0 atm = 101.3 kPa = 1.013 x 105 Pa)
Understand that you must use kPa (pressure) and dm3 or L (volume) when using R = 8.314 J/mol K
Use the ideal gas equation to calculate the molar mass of a gas, given the necessary information?
Gas Stoichiometry Practice
For all of these problems assume that the reactions are being performed at a pressure of 100 kPa
and a temperature of 273 K and so the molar gas volume is 22.7 dm3.
1)
Calcium carbonate decomposes at high temperatures to form carbon dioxide and calcium
oxide:
CaCO3(s) → CO2(g) + CaO(s)
How many grams of calcium carbonate will I need to form 3.45 dm3 of carbon dioxide?
2)
Ethylene burns in oxygen to form carbon dioxide and water vapor:
C2H4(g) + 3 O2(g) → 2 CO2(g) + 2 H2O(g)
How many liters of water can be formed if 1.25 dm3 of ethylene are consumed in this reaction?
3)
When chlorine is added to acetylene, 1,1,2,2-tetrachloroethane is formed:
2 Cl2(g) + C2H2(g) → C2H2Cl4(l)
How many dm3 of chlorine will be needed to make 75.0 grams of C2H2Cl4?
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Additional Questions - Combined and Ideal Gas Equations
1. The pressure on 600cm3 of gas is increased from 100kPa to 300kPa at constant temperature.
What will the new volume of gas be?
2. 1dm3 of gas in a container at -73oC is allowed to expand to 1.5dm3. What must the
temperature be so that the pressure remains constant?
3. 4.00dm3 of air at 0oC and a pressure of 2.00 atmospheres, is heated to 273oC and the
pressure is increased to 8.00 atmospheres. What will the new volume of gas be?
4. A flask contains 250cm3 of gas at 65oC and atmospheric pressure. The flask is then heated to
650oC. What factor does the pressure increase be?
5. A syringe contains 50cm3 of gas at 1.0 atm and 20oC. What would the volume be if the gas
were heated to 100oC, at the same time compressing it to 5.0atm?
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6. 3.376g of a gas occupies 2.368 dm3 at 17.6oC and a pressure of 96.73kPa, what is its molar
mass?
7. What volume is occupied by 0.0200g of oxygen gas at 27oC and a pressure of 107kPa?
8. A steel cylinder contains 32dm3 of hydrogen at 4.9 x 106 Pa and 39oC. Calculate:
a. The volume that the hydrogen would occupy at s.t.p. (0oC and 101.3kPa)
b. The mass of hydrogen in the cylinder.
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Lesson 17: Unit 2 Lesson 6 - Finding Molar Mass of a Gas
[Top of Doc]
Learning Goals:
PROBLEM-SOLVING
Can you use the ideal gas equation to calculate the molar mass of a gas, given the necessary
information?
EXPERIMENTAL SKILLS
Can you create and/or use the correct data table? This includes correct labels, units and consistent use
of decimal places.
Can you gather and record data in a safe and accurate manner?
Can you analyse your data and communicate it in an appropriate manner?
Can you uses critical thinking skills to evaluate information and drawing appropriate conclusions?
Can you evaluate weaknesses and errors in your method based on experimental evidence?
Finding Molar Mass of a Gas Laboratory
By now you have been introduced to the Ideal Gas Law, PV = nRT, where
P= pressure (kPa)
V= volume (dm3)
n= number of moles
R= the Ideal Gas Constant (R = 8.31 J mol-1 K-1 or 0.0821 L atm mol-1 K-1)
T = temperature (K)
This equation can be used to determine the molar mass of an ideal gas. All you need to do is:
1. Know the mass of the gas (the “g” in g/mol). You can normally get that from a balance.
2. Know the moles (n) of the gas (the “mol” in g/mol). You get that from the ideal gas equation.
So, to determine the molar mass of an unknown gas, what data would you need to collect and what
instruments would you use to collect this data? Write your response below.
In this laboratory, you will make measurements to enable you to calculate the Mass of the gas in a simple
lighter. Pay close attention to your units!
Procedure
1) Dip a lighter into the water, dry it very carefully, then weigh it and record its mass.
2) Obtain a tray of water and a eudiometer. Fill the eudiometer completely with water and carefully invert
it into the tray of water so that NO AIR is in the eudiometer. Carefully invert a funnel into the opening of
the eudiometer. Again do not allow any air to enter the eudiometer.
3) Measure the temperature of the water and record.
4) Place the lighter under the funnel and press the release valve so that the gas exits the lighter and
passes through the funnel into the eudiometer.
5) Collect enough gas so that the level of the gas inside the eudiometer is the same as the water level
outside of it.
6) Record the volume of the gas collected.
7) Record the atmospheric pressure (remember that it must be in kPa!).
8) Carefully dry the lighter, being sure to remove all water from the exterior and interior.
9) Reweigh the lighter and determine the amount of gas released.
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(Support Video : Finding MM of a Gas - video of procedure (6 minutes)
Data Table
Measurement
Value
Uncertainty
Initial mass of lighter (g)
Final mass of lighter (g)
Volume of gas (cm3)
Temperature of water (oC)
Temperature of water (K)
Atmospheric pressure (kPa or
atm…depending on your R)
Calculations
Use the data that you have collected to calculate the molecular mass of the lighter’s gas. You can check your
calculations with the answer checker on the website (Make your own copy).
Conclusion and evaluation
1) The gas was butane (C4H10). Use this information to calculate the percent error of your measured
value.
2) Which measurement(s) do you feel had the greatest uncertainty?
3) Briefly evaluate the procedure. What systematic errors did you encounter? Don’t forget the direction of
the errors.
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4) Suggest how the procedure could be improved.
5) Complete the following table based on your observations and processed data values (Add
additional rows if needed).
Weakness
Explanation of Weakness
based on
Results/Observations
Improvements
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Lesson 18: Unit 2 Lesson 7 - Intro to Solutions and Concentration
[Top of Doc]
Learning Goals:
Understanding
Can you describe the meanings of the words mixture, solute, solvent and solution?
Can you define concentration of solutions using molarity (moles of solvent/dm3 of solution)?
EXPERIMENTAL SKILLS
Can you express solution concentration in mol/dm3?
Can you express solution concentration in g/dm3?
Calculate volumes of solutions given mass or moles?
Can you calculate the number of ions in a solution?
Calculate moles of solutions given volume and/or mass?
Part A: Mass and Volume Student Laboratory Activity
Introduction
A solution is a homogenous mixture, a mixture that has uniform composition and properties throughout.
Aim: To investigate solutions where different numbers of moles of nickel (II) sulphate are dissolved in various
volumes of water.
Materials Per Group
Weighing Boat
50 ml Beaker
Glass Stirring Rod
50 ml Graduated
Cylinder
Cuvette with Lid
Procedure
1.
2.
3.
4.
Weigh out 1.577 g of nickel sulphate in a 50 cm3 beaker.
Add 10 cm3 of water and stir with a glass rod to dissolve the solid.
Pour the solution into a measuring cylinder and top up the solution to 15 cm3.
Repeat steps 1-3 using the amounts stated in the table below so that you have a total of 5 solutions.
Mass (g)
#
+/-
Total Volume
(cm3)
+/-
1
1.577
15
2
2.103
20
3
2.629
25
Measured Concentration
(mol / dm3)
+/-
75
4
3.154
30
5
3.680
35
You are now going to test your solutions in the spectrophotometer. This instrument measures how many
moles there are in a coloured solution by detecting how much light passes through the solution. It then
calculates how many moles would be present in a dm3, which is the number that you will read.
5. Holding by the ridged sides only, wash a cuvette out with distilled water and then with a small amount
of the solution that you wish to measure (in this case, the first solution, which has a volume of 15 cm3).
6. Fill the cuvette with the 15 cm3 solution until it is ¾ full. Put a lid onto the cuvette and wipe the outside
with soft tissue so that it is clean and dry.
7. Place the cuvette into the spectrophotometer with the smooth side facing the arrow. Close the lid and
record the reading in the table above.
8. Repeat steps 5 – 7 with the other solutions that you have made.
Results and calculations
Note your observations about the five solutions below.
Complete the table below using the values of mass and volume for the solutions that you made. Note: For
molar mass - nickel (II) sulphate hexahydrate (NiSO4.6H2O)
Moles in 1 dm3 of
Mass
Moles
Volume
Volume (dm3)
solution
(cm3)
Check your answers using: 01 Moles and Volume Lab Checker
Conclusion
What relationship do you find between the mass (and therefore moles) and volume of the solutions that you
made and the concentrations that you read from the spectrophotometer?
……………………………………………………………………………………………………………………………...
………………………………………………………………………………………………………………………………
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Part B Calculations Involving Concentration
1. How many moles of hydrochloric acid are present in 0.80 dm3 of a solution with a concentration of 0.4
mol dm-3?
2. Sodium phosphate has the formula Na3PO4. What is the concentration of sodium ions in a 0.6 mol dm-3
solution of sodium phosphate?
3. What volume of a 0.5 mol dm-3 solution of sodium hydroxide can be prepared from 2 g of the solid?
4. What are the concentrations of the solutions produced by dissolving:
a. 3.0 moles of nitric acid in 4.0 dm3 of solution?
b. 2.81g of KOH in 2.00 dm3 of solution?
c. 5.00g of magnesium sulfate heptahydrate in 250 cm3 of solution?
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5. How many moles are in the following?
a. 7.0 dm-3 of sulphuric acid of concentration 0.30 mol dm-3.
b. 50cm3 of a 0.040 mol dm-3 solution of lithium chloride.
c. 15.0 cm3 of a solution made by dissolving 5.80g of zinc chloride in 2.50 dm3 of solution.
6. What volume of solution could you produce in the following cases?
a. 1 mol dm-3 copper (II) chloride from 0.4 moles of the solid.
b. 0.0200 mol dm-3 NaNO3 starting from 5.00g of the solid.
c. Dilute 0.50 mol dm-3 sulphuric acid starting with 20 cm3 of a concentrated 18 mol dm-3 solution.
(note if you’re not sure how to do this, don’t worry we will discuss dilutions in detail next lesson)
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7. Describe how you would prepare 500 cm3 of a 0.100 mol dm-3 NaCl solution?
8. How would you prepare 1.2 dm3 of a 0.40 mol dm-3 solution of HCl starting from a 2.0 mol dm-3
solution? (note if you’re not sure how to do this, don’t worry we will discuss dilutions in detail next
lesson)
9. 500 cm3 of 0.500 mol dm-3 NaCl is added to 500 cm3 of 1.00 mol dm-3 Na2CO3 solution. Calculate the
final concentration of Na+ ions in solution.
10. CHALLENGE: When hydrochloric acid is added to aqueous lead nitrate, solid lead chloride is
precipitated. If 10 cm3 of 2 mol dm-3 hydrochloric acid is added to 40 cm3 of 0.5 mol dm-3 aqueous lead
nitrate…what is the concentration in the final solution of:
a. What is the balanced equation?
b. What is the limiting reagent?
c. What is the concentration in the final solution of:
1. nitrate ions?
2. chloride ions
3. hydrogen ions
4. lead ions
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Lesson 19: Unit 2 Lesson 8 - Making Standard Solutions and Dilutions
[Top of Doc]
EXPLAINING
Can you explain how to dilute a solution using appropriate lab materials?
Can you explain how to create a standard solution using appropriate lab materials?
PROBLEM SOLVING
Calculate moles of solutions given volume and/or mass?
Can you express solution concentration in mol/dm3?
Can you determine the volumes of solutions needed to create dilutions using the dilution equation, CCVC
= CDVD?
EXPERIMENTAL SKILLS
Can you create a standard solution using appropriate lab materials?
Can you dilute a solution using appropriate lab materials?
What is a standard solution?
A standard solution is a solution whose concentration is known accurately. Its concentration is usually given
in mol dm–3. When making up a standard solution it is important that the correct mass of substance is
accurately measured. It is also important that all of this is successfully transferred to the volumetric flask used
to make up the solution. The following procedure will make sure that this happens.
Background calculations
1. Work out the number of moles needed to make up
a solution with the required volume and
concentration.
2. Now work out the relative formula mass, Mr, of the
chosen substance.
3. Work out the mass of substance needed using
your answers from steps 1 and 2.
Making up the solution
• Take a weighing dish and weigh out the required mass of substance (A). Transfer this amount to a beaker
then reweigh the weighing dish (B). Subtract this mass from the first mass (A-B) to find the mass of
substance added to the beaker.
• Transfer this amount to a beaker. Add the minimum amount of water from a wash bottle to dissolve it.
• Stir with a glass rod until all the solid is dissolved, and then transfer the solution to the volumetric flask.
Use more water from the wash bottle to rinse out the beaker and the glass rod. Do this at least twice.
• Add water to just below the line on the volumetric flask. Add the final drops with a teat pipette to ensure
that the bottom of the meniscus is on the line.
• Put the lid on the flask and turn the flask over a couple of times to mix the solution.
Homogeneous mixtures
Solutions are homogeneous mixtures. For example, if we have a beaker with distilled water and dissolve 3.0 g
of NaCl, then all of the NaCl particles will be distributed evenly throughout the whole solution. This means that
if we take any volume of the solution, the number of particles of NaCl dissolved per unit of volume will be
exactly the same.
3
Let’s say that the solution has a total volume of 1 L, which equals 1 dm . If we take a 500 cm3 sample of the
solution, then it will contain 1.5 g of NaCl.
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Concentration (Molarity):
Concentration(M) =
number of moles of solute (n)
Volume of solution (expressed always in Liters)
*Note: remember that the solution involves the solute + the solvent
i.e. A 1.0 M solution of NaCl can be prepared by obtaining 1 mol of NaCl (58.5g), placing it in a volumetric
flask, and adding enough water until the 1.00L mark is reached. Notice that you do not add 1.00 L of water,
but add enough water to the solute, until the 1.00 L mark is reached*.
Making a Standard Solutions Laboratory Activity
Part A: Making a standard solution from a solid salt.
Your teacher will assign you one of the following solutions to prepare.
Volume
(mL)
Concentration
(M)
1
50.00
0.054
2
50.00
0.036
3
50.00
0.096
4
50.00
0.088
5
50.00
0.063
Solution #
Mass of CuSO4*5H2O (g)
02 Making Standard Solutions Lab Checker
Calculate the mass of CuSO4*5H2O required for your solution in the space below.
Part B: Making a standard solution by diluting your stock solution.
Use the solution you created in Part A to create a new 25.00 mL solution (Solution B) whose concentration is
1/5 of the concentration in Part A.
Solution
#
Volume of Solution B
(V2, mL)
Concentration of
Solution B
(M2, M)
Volume of Solution A
(V1, mL)
Concentration of
Solution A
(M1, M)
(Given Part A Table)
25.00
02 Making Standard Solutions Lab Checker
Complete the calculations required for your dilution in the space below.
Step 1: Determine M2
Step 2: Determine V1, use M1V1= M2V2
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Making Standard Solutions Calculation Practice
Remember that you can change the concentration of a solution by adding more solvent. While you
cannot increase the concentration of a solution in this manner, you can create a more dilute solution
by increasing the amount of solvent. You can determine the amount of a solution needed to dilute by
using the following:
M1 x V1 = M2 x V2
Where M = molarity and V = volume. M1 and V1 are the initial solution’s molarity and volume, while
M2 and V2 are the final solution’s molarity and volume. If needed, you can find the molarity of a
solution by the usual formula: M = moles of solute/ liters of solution
Use the formula and information above to solve the following problems. Show your work and watch
labels. The first three problems are questions regarding molarity and the others involve the dilutions
formula above.
1. Determine the number of grams of NaHCO3 that are in one liter of a 2.1 M solution
2. Determine the number of grams of NaNO3 that are present in 500 mL of a 1.0 M solution.
3. Determine the number of grams of CCl4 in 450 mL of a 3.2 M solution.
4. You need to make 300 mL of a 0.40 M solution of sodium chloride. The only available solution
is 1.0 M. Determine how to make the needed dilution.
5. You have to make 500 mL of a 0.50 M BaCl2 . You have 2.0 M barium chloride solution available.
Determine how to make the needed dilution
82
6. You need to make 10.0 L of 1.2 M KNO3 . What molarity would the potassium nitrate solution need
to be if you were to use only 2.5 L of it?
7. Using a 4.0 M solution of MgSO4 , determine how to make 300 mL of a 1.7 M dilution.
8. If you dilute 174 mL of a 1.6 M solution of LiCl to 1.0 L, determine the new concentration of the
solution.
9. One liter of a solution is prepared by dissolving 125.6 g of NaF in it. If I took 180. mL of that
solution and diluted it to 500. mL, determine the molarity of the resulting solution.
10. Challenge Question: Exactly 16.0 mL of a solution A is diluted to 300 mL, resulting in a new
solution B that has 0.50 M concentration. If the solution was made with NaCl, determine the number
of grams of NaCl needed to make 1.00 L of the original solution A.
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Lesson 20: Unit 2 Lesson 9 - Beer-Lambert Law Theory and Demo
[Top of Doc]
UNDERSTANDING
Do you understand that spectroscopy can determine the concentration of a solution?
Do you understand that Beer-Lambert law relates the absorption of light by a solution to three variables according to the
equation: A = abc, where a = molar absorptivity (how intensely a sample of molecules or ions absorbs light of a specific
wavelength), b The path length, and c = the concentration of the solution?
Do you understand that in experiments the path length and wavelength of light are held constant, making the absorbance
proportional only to the concentration of absorbing molecules or ions?
EXPLAINING
Can you explain the relationship between the light absorbed by a solution of molecules or ions in relationship to the
concentration, path length, and molar absorptivity?
PROBLEM SOLVING
Can you solve algebraic and graphical problems involving Beer-Lambert Law?
EXPERIMENTAL SKILLS
Can you use a SpectraVis or Colorimeter to determine the absorption and/or concentration of a solution?
Beer-Lambert Law - Worksheet
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Lesson 21: Unit 2 Lesson 10 - Introduction to Titrations
[Top of
Doc]
UNDERSTANDING
Can you calculate moles or concentration of substance in a titration?
EXPERIMENTAL SKILLS
Can you conduct a titration using appropriate lab materials?
How to use it
1. Fix the burette into the burette holder, taking care that it is vertical and stable. Place a beaker
underneath the burette.
2. Close the tap, and run some deionized water into the top of the burette. Let the water clean the inside
of the burette. Open the tap, and allow the water to drain out. Repeat.
3. Close the tap, and (using the funnel) run some of the required reagent, e.g. acid, into the top of the
burette. Open the tap, and allow the reagent to drain through into the beaker. Repeat.
4. Close the tap, and fill the burette to just above the 0.00 cm3 mark with the required reagent. Remove
the funnel. Make sure that there are no air bubbles inside the burette. Slowly open the tap, and allow
the reagent to run down to (or just past) the 0.00 cm3 mark. Close the tap.
5. Remove the beaker, and place a white tile under the burette. Put a conical flask under the burette, and
adjust the height of the burette so that the tip is just above the lip of the conical flask. The burette is
now ready for use.
Recording the results
1. Construct Results tables like the ones below. Before you start, record the reagents used.
Here is a video for students who are unable to watch the live demonstration in class. Setting up
and performing a titration (6:52)
Burette reagent and concentration
Pipette reagent
(plus, record the volume and uncertainty)
Indicator
Rough
Trials
1
2
3
4
5
Final volume (cm3 ± 0.05cm3)
Initial volume (cm3 ± 0.05cm3)
Titre (cm3 ± 0.10cm3)
Mean titre (cm3)
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2. In the first run, open the burette tap and allow the liquid to run into the conical flask, swirling all the
time. Close the tap as soon as the indicator changes color. You should overshoot the end-point a little.
Record the start and end volumes as you go. Record volumes to the nearest 0.05cm3, i.e. all volumes
should end in .x0cm3 or .x5cm3.
3. In subsequent runs, allow the burette reagent to run through until approximately 1cm3 before the
end-point found in the rough run. Close the tap until the burette is letting out one drop at a time. As
soon as the indicator changes color, close the tap completely.
4. If your three accurate runs are very different from each other, repeat until a more consistent result is
obtained, i.e. to concordance (± 0.10cm3). Put a pencil tick against the titres you use in your calculation
of the mean titre.
5. Calculate the mean volume delivered (the titre) and record it in the table. Concordant titres must be
used to calculate the mean. Concordant titres is generally taken to mean a minimum of 2 titres that
are within 0.1 cm3 of one another. For example, 21.45 cm3 and 21.55 cm3 are concordant but 21.45
and 21.60 cm3 are not concordant.
HOMEWORK ASSIGNMENT FOR LESSON 3 THE BURETTE
DEMONSTRATION
Use the following data to determine the unknown concentration of HCl.
1. Mean titre calculation:
2. Unknown concentration of HCl calculation:
88
Lesson 22: Unit 2 Lesson 11
Formative Lab 2 Standardization and Titration
[Top of Doc]
Standardize Hydrochloric Acid Using Sodium Carbonate Solution
Part 1: Preparing a standard solution of sodium carbonate
Introduction
A standard solution is one whose concentration is known exactly. Standard solutions of liquids, for example
acids, are easy to prepare and are usually supplied. Standard solutions of solids can be prepared by
weighing a mass of solid, and dissolving it in a known volume of solution in a volumetric flask. Today, you are
going to prepare a standard solution of sodium carbonate which you will then use to standardize hydrochloric
acid.
Procedure Part 1:
1.
Weigh a weighing boat then weigh approximately between 1.2g and 1.4g of sodium carbonate into it.
2.
Transfer the contents a 250cm3 beaker. Weigh the weighing boat again.
The difference between the two accurate masses is the mass of sodium carbonate in your beaker.
3.
Add distilled water cautiously down the side of the beaker. Use the minimum amount of water to
dissolve the solid, and swirl the beaker to mix the contents. Stir using a glass rod to dissolve the solid
completely.
4.
Transfer the solution into a 250 cm3 volumetric flask. Wash the beaker, rod and funnel several times
using distilled water from the wash bottle, letting the washings go into the flask.
5.
Make up to the mark on the volumetric flask with distilled water. Stopper firmly, and shake the flask
thoroughly to mix the contents.
Data Table Part 1—Sodium Carbonate
Read the procedure carefully and create the necessary data table for ALL the necessary data (e.g. how many
weighings will be required?).
Create a Data Table Below:
89
Part 2: Standardizing Hydrochloric acid
Introduction
You will now find the concentration of dilute hydrochloric acid by titration. The reaction between sodium
carbonate and hydrochloric acid takes place as follows
Na2CO3(aq) + 2HCl(aq) → 2NaCl(aq) + CO2(g) + H2O(l)
Procedure Part 2:
1.
Using a volumetric pipette, transfer a 20 cm3 aliquot (portion) of your sodium carbonate solution to a
250cm3 capacity conical flask. Add a few drops of methyl orange indicator solution.
2.
Titrate with the unknown concentration of hydrochloric acid. The end-point of the titration is when the
solution just changes from yellow to red.
3.
Repeat steps 1 - 2 until concordance (i.e. until the readings are the same or within 0.1cm3).
NOTE: This titration is a bit different from a normal titration. Normally the burette contains the reagent with a
known concentration and the pipette contains the reagent with the known concentration. In this lab it is the
other way around.
Pipette Reagent: Na2CO3 (KNOWN CONCENTRATION)
Burette Reagent: HCl (UNKNOWN CONCENTRATION)
Data Table Part 2--Titration
Create a set of data tables based on The Burette demonstration. Don’t forget to record qualitative data as
well.
Sample Data (if you have missed the class)
90
Calculations 04 Standardise HCl Using Na2CO3 Lab Checker
Record all of your answers to 4 decimal places (or 4 significant figures).
1.
Calculate the number of moles of sodium carbonate in the mass used.
2.
Calculate the concentration of the sodium carbonate primary standard.
3.
Calculate the moles of sodium carbonate in the 20 cm3 aliquots that you titrated.
4.
How many moles of HCl will react with that number of moles of Na2CO3?
5.
Find the concentration of HCl using your mean titre.
6.
If the theoretical concentration of HCl is 0.1000 M, calculate the percent error in your experimental
value.
91
Lesson 24: Unit 2 Lesson 13 - Formative Quiz 2/Unit 2 Review
[Top of Doc]
Gases Review Questions
1. A compound has the empirical formula CHCl. A 256 cm3 flask at 100 oC and 100 kPa contains
0.800 g of the gaseous compound.
a. Calculate the molar mass of the compound.
b. Determine the molecular formula of the compound.
2. An unknown diatomic gas with a volume of 2.5 dm3 has a mass of 7.91 g at 273 K and 1.01 x 105
Pa.
a. Calculate the molar mass of the gas.
b. What is the identity of the gas?
3. A chemist has synthesized a greenish-yellow gaseous compound of chlorine and oxygen and
finds that is density is 7.71 g/dm3 at 36 oC and 2.88 atm. Calculate the molar mass of the
compound.
4. At 35oC, 0.210 L of a compound exerted a pressure 172.2 kPa. If the mass of the compound was
2.38 g, what was the molecular mass of the compound?
5. Sodium azide (NaN3) is used in some automobile air bags. The impact of a collision triggers the
decomposition of NaN3 as follows:
NaN3 (s) → Na (s) + N2 (g)
The nitrogen gas produced quickly inflates the bag between the driver and the windshield and
dashboard. Calculate the volume of N2 generated at 80 oC and 1.08 atm by the decomposition of
60.0 g of sodium azide.
92
6. Hydrochloric acid reacts with solid magnesium metal to produce a solution of magnesium chloride
and hydrogen gas.
a. Write the equation for the reaction that occurs.
b. What volume of hydrogen gas can be generated from 50.0 mL of 2.0 mol/dm3 hydrochloric
acid? Assume that the reaction occurs at STP and the molar volume of gas is 23 dm3/mol.
7. Methane, the principal component of natural gas, is used for heating and cooking.
a. Write an equation for the combustion of methane, CH4.
b. If 15.0 moles of methane are reacted, what is the volume of carbon dioxide (in liters)
produced at 23.0 oC and 0.985 atm?
8. In alcohol fermentation yeast converts glucose to ethanol and carbon dioxide:
C6H12O6 (s) → C2H5OH (l) + CO2 (g)
a. If 5.97 g of glucose are reacted, what volume of carbon dioxide is expected at STP? Assume
a molar volume of 23 dm3/mol (for this specific question).
b. If 1.25 dm3 of carbon dioxide are measured, what is the percent yield of the reaction?
93
9. Calculate the volume from the given number of moles of gas. 1 mole of gas occupies a volume of
22.4 dm3 at 293 K and 1 atm pressure.
a. 1 mole of CO2
b. 3.4 moles of HBr
c. 0.1 moles of NH3
d. 0.11 moles of Cl2
10. Calculate the number of moles from the given volume of gas. (Assume 22.7 dm3 can be used)
a. 200 cm3 of CO2
b. 226 cm3 of HBr
c. 500 cm3 of NH3
d. 256 cm3 of Cl2
11. Calculate the molar mass of a gas from volume and mass data. (Assume 22.7 dm3 can be used)
a. 0.467 g of gas occupies 131 cm3
b. 0.218 g of gas occupies 84 cm3
c. 0.296 g of gas occupies 93 cm3
d. 0.267 g of gas occupies 187 cm3
94
Solution Review Questions
1.
A student carried out an experiment to determine the concentration of ethanoic acid in a
solution of vinegar.
•
The student used a measuring cylinder to measure out 25.0 cm3 of the vinegar solution.
•
This solution was then transferred to a 250 cm3 volumetric flask and the liquid level was
carefully made up to the mark with distilled water.
•
A pipette was used to transfer 25.0 cm3 portions of the acidic solution to conical flasks.
•
The solution was then titrated with sodium hydroxide solution, concentration 0.100 mol
dm–3, using phenolphthalein as the indicator.
CH3COOH(aq) + NaOH(aq) → CH3COONa(aq) + H2O(l)
Results
Titration number
1
2
3
4
Burette reading (fi nal) / cm3
28.55
28.00
40.35
28.05
Burette reading (initial) / cm3
0.00
0.05
12.30
0.05
Volume of NaOH used / cm3
28.55
27.95
28.05
28.00
(a)
In this titration, what is the colour change of the phenolphthalein indicator?
From .......................................................................... to
..............................................................................
(2)
(b)
Explain why the mean titre should be based only on titrations 2, 3 and 4.
..............................................................................................................................................
..............................................................................................................................................
(1)
(c) Calculate the mean titre in cm3.
(1)
(d) (i)
Using your answer to (c), calculate the number of moles of sodium hydroxide in the
mean titre.
(1)
(ii)
Hence state the number of moles of ethanoic acid, CH3COOH, in 25.0 cm3 of the
diluted solution used in the titration.
(1)
95
(iii) Calculate the concentration of the diluted acid solution in mol dm–3.
(1)
(iv) Hence calculate the concentration of the ethanoic acid in the original vinegar
solution in mol dm–3.
(1)
(v)
Use your answer from (d)(iv) to state the concentration of the ethanoic acid in the
original vinegar solution in units of g dm–3.
[The molar mass of the ethanoic acid is 60 g mol–1.]
(1)
(e)
Suggest, with a reason, how the student’s method of preparing the diluted solution could
be improved.
Improvement
..............................................................................................................................................
Reason
..............................................................................................................................................
(2)
(f)
The burette used in the titration had an uncertainty for each reading of
(i)
0.05 cm3.
Identify which ONE of the following should be regarded as the true value of the titre
in titration number 2?
A
Between 27.90 and 28.00 cm3
B
Between 27.925 and 27.975 cm3
C
Between 27.85 and 28.05 cm3
(1)
(ii)
Suggest ONE reason why a student may obtain volumes outside the uncertainty of
the burette when performing a titration.
.....................................................................................................................................
.............................................................................................................................(1)
96
2.
Wine is an aqueous solution of ethanol with traces of other organic compounds which give the
wine its characteristic flavour and aroma. Once opened, oxidation of the ethanol in the wine
produces ethanoic acid.
• A white wine with an ethanol concentration of 2.25 mol dm–3 was opened and allowed to
stand at room temperature for 2 weeks.
• A 25.0 cm3 sample of the wine was transferred to a clean conical flask and phenolphthalein
indicator added.
• Aqueous sodium hydroxide of concentration 0.205 mol dm–3 was added from a burette until
the colour of the indicator changed.
• The titration was repeated and the titre values in cm3 were 26.35, 26.90 and 26.45.
The equation for the neutralisation reaction is
CH3COOH + NaOH → CH3COONa + H2O
(a)
(i)
Name the piece of apparatus used to measure 25.0 cm3 of wine.
.....................................................................................................................................
(1)
(ii) State how the burette should be rinsed.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(1)
(iii) State the colour change at the end-point.
From ...................................................................... to
........................................................................
(2)
(b) (i)
What is meant by the term concordant results?
.....................................................................................................................................
(1)
(ii) Calculate the mean (average) titre which will be used to calculate the concentration
of ethanoic acid.
(1)
(iii) Calculate the number of moles of sodium hydroxide reacting with 25.0 cm3 of the
wine. (1)
97
(iv) Hence calculate the concentration of the ethanoic acid, in mol dm–3.
(2)
(v)
Calculate the percentage of the ethanol that has oxidised, given that one mole of
ethanol forms one mole of ethanoic acid.
(1)
(c)
Suggest why this method would not be effective for the analysis of the acid content of a
red wine.
..............................................................................................................................................
..............................................................................................................................................
(1)
3. A titration is carried out by adding sodium hydroxide solution from a burette to 25.0 cm3 of
aqueous 0.0500 mol dm–3 butanedioic acid, (CH2COOH)2, to which a few drops of phenolphthalein
have been added.
(CH2COOH)2(aq) + 2NaOH(aq) → (CH2COONa)2(aq) + 2H2O(l)
Diagram I
(a)
The burette readings recorded by a student carrying out the titrations are shown in the
table below.
Titration numbers
1
2
3
Burette reading
(final) / cm3
23.90
23.60
23.65
Burette reading
(initial) / cm3
0.00
0.00
0.15
Titre/cm3
23.90
23.60
23.50
Used in mean (
)
98
(i)
On Diagram II below, show the level of the sodium hydroxide solution when the
final burette reading is recorded in titration 3.
Diagram II
(1)
(ii) Calculate the mean (or average) titre.
Show which titres you have used in your calculation by putting a tick (
the table above. (2)
) in the appropriate boxes in
(b)(i)Calculate the amount (moles) of butanedioic acid, (CH2COOH)2, in 25.0 cm3 of the 0.0500 mol
dm–3 solution.(1)
(ii) Calculate the amount (moles) of sodium hydroxide, NaOH, in the mean titre.(1)
(iii) Calculate the concentration of the sodium hydroxide solution in mol dm–3. (1)
99
4.
(a)
Calculate the concentration, in mol dm–3, of a solution containing 1.28 g of anhydrous
sodium carbonate, Na2CO3, in 250 cm3 of solution
(3)
(b)
In a series of titrations, hydrochloric acid was added, from a burette, to 25.0 cm3 portions
of the sodium carbonate solution pipetted into conical flasks. Methyl orange was added
as the indicator.
The burette readings are shown in the table below.
1
2
3
Burette reading at
end/cm3
24.80
48.90
24.40
Burette reading at
start/cm3
0.00
24.80
0.00
Titre/cm3
24.80
24.10
24.40
Number of titrations used to calculate the mean (average) titre: 1, 2 and 3
Mean titre = 24.43 cm3 of hydrochloric acid
(i)
Give the colour change that would be observed at the end point.
From ................................................. to ......................................................
(1)
(ii)
The student carrying out the titrations was criticised by the teacher for not carrying
out at least one more titration.
Suggest a reason why the teacher’s criticism was justified.
.....................................................................................................................................
.....................................................................................................................................
(1)
(c) Using the mean titre given and your answer to (a)(ii), calculate the concentration of the
hydrochloric acid in mol dm–3. The equation for the reaction in the titration is:
Na2CO3(aq) + 2HCl(aq) → 2NaCl(aq) + CO2(g) + H2O(l)
100
(2)
(d) Before titration 2, the student rinsed the pipette with water and then immediately used it
to transfer sodium carbonate solution to the conical flask for the titration.
If 0.5 cm3 of water was present in the pipette, what affect would this have on the titration
result.
……………………………………………………………………………………………………………
…………………………………………………………………………………………………………
(1)
5.
(a)
The reaction involved in a titration is
(COOH)2(aq) + 2NaOH(aq) → (COONa)2(aq) + 2H2O(l)
25.0 cm3 of the aqueous 0.0500 mol dm–3 ethanedioic acid required 25.50 cm3 of the
aqueous sodium hydroxide for neutralisation.
(i)
Calculate the amount (moles) of ethanedioic acid in 25.0 cm3 of the solution.
(1)
(ii)
Calculate the amount (moles) of sodium hydroxide in 25.50 cm3 of the solution.
(1)
(iii) Calculate the concentration of the sodium hydroxide solution in mol dm–3.
(1)
(c) Calculate the mass of hydrated ethanedioic acid, (COOH)2.2H2O, needed to make up
200 cm3 of aqueous 0.0500 mol dm–3 solution.
(3)
101
(d)
When making up the solution of ethanedioic acid a student, by mistake, uses a 200 cm3
instead of a 250 cm3 volumetric flask. The student dissolves the mass of ethanedioic acid
crystals calculated to make up 250 cm3.
Explain what effect this would have on the student’s volume of sodium hydroxide solution
used in the titration.
[No calculation is required in your answer.]
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
(2)
6.
A 1.62 g sample of impure sodium carbonate was dissolved in distilled water and then made
up to 250 cm3. 25.0 cm3 of this solution was put into a conical flask and three drops of methyl
orange indicator added. This was titrated against a 0.105 mol dm–3 solution of hydrochloric
acid until the end point was reached. The titration was repeated three more times. The results
are shown below.
1
2
3
4
Burette reading (final)
25.30
25.30
25.85
25.95
Burette reading (at
start)
0.00
0.50
0.75
1.25
Titre/cm3
25.30
24.80
25.10
24.70
The equation for the reaction is:
Na2CO3 + 2HCl → 2NaCl + H2O + CO2
(a)
The student was supplied with a burette that may not have been clean. What
precautions should be taken before filling it with the standard hydrochloric acid solution?
.............................................................................................................................................
.............................................................................................................................................
.............................................................................................................................................
(2)
(b) (i)
Select the appropriate titres and calculate their mean.
..............................................................................................................................................
..............................................................................................................................................
(2)
102
(ii)
Calculate the amount (in moles) of hydrochloric acid solution in the mean titre.
(1)
(iii) Calculate the amount (in moles) of pure sodium carbonate in 25.0 cm3 of solution.
(1)
(iv) Calculate the amount (in moles) of pure sodium carbonate in 250 cm3 of solution.
(1)
(v)
Calculate the mass of pure sodium carbonate, Na2CO3, taken. (2)
(vi) Calculate the percentage purity of the sample of sodium carbonate. (1)
7. A laboratory technician is given the task of making up 5 dm3 of aqueous sodium hydroxide of
concentration 0.100 mol dm–3. The technician finds the following data on sodium hydroxide.
Formula NaOH
Soluble in water
Solid which absorbs moisture and acidic gases from the air
Solid is corrosive
Reacts with acids in aqueous solution
e.g. 2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(i)
103
The technician prepares the solution and checks its concentration, following the procedure outlined
below.
(a)
I
The technician calculates the mass of sodium hydroxide needed to make 5 dm3 of
0.100 mol dm–3 solution.
II
The technician adds 5 dm3 of water to a plastic bucket.
III
The technician weighs the calculated mass of sodium hydroxide, transfers it to the
plastic bucket and stirs until the sodium hydroxide has dissolved.
IV
The technician titrates 25.0 cm3 samples of the sodium hydroxide solution with
0.0500 mol dm–3 sulphuric acid.
V
The mean titre is 23.50 cm3 of 0.0500 mol dm–3 sulphuric acid.
Calculate the mass of sodium hydroxide that the technician needs to take, to make 5 dm3
of solution of concentration 0.100 mol dm–3.
(2)
(b)
Calculate the concentration, in mol dm3, of the sodium hydroxide solution from the
titration results in IV and V.
(3)
(c) The actual concentration of the sodium hydroxide solution is not exactly 0.100 mol dm–3
as the technician intended.
(i)
Suggest ONE reason for this, which is a consequence of the way in which the
technician makes up the solution.
.....................................................................................................................................
.....................................................................................................................................
(1)
(ii) Suggest ONE reason for this, which is a consequence of the chemical properties of
the sodium hydroxide.
.....................................................................................................................................
.....................................................................................................................................
(1)
104
Unit 3: Bonding and Electrons - Learning Goal Summary
Can You? / Do You?
UNDERSTANDING
Understand that equal volumes of gases, under the same conditions of temperature and pressure,
contain the same number of particles? (Avogadro’s Law)
Understand the four assumptions of kinetic molecular theory (KMT)?
EXPLAINING
Explain gas phenomena (pressure, temperature, volume conditions or changes) using the ideas of
kinetic molecular theory?
PROBLEM SOLVING
Solve problems related to Avogadro's Law (using volumes of gases and stoichiometric analysis,
including the concept of limiting reagents)?
Convert Celsius and Kelvin temperatures?
Use the combined gas law to solve problems concerning pressure, volume and temperature?
Use the ideal gas law to solve problems concerning pressure, volume, temperature and moles of gas?
Convert between pressure units? (1.0 atm = 101.3 kPa = 1.013 x 105 Pa)
Use the molar gas volume, taken as 22.7 dm3 at room temperature and pressure.
Understand that you must use kPa (pressure) and dm3 or L (volume) when using R = 8.314 J/mol K
Use the ideal gas equation to calculate the molar mass of a gas, given the necessary information?
EXPERIMENTAL SKILLS
Write a research question, identify relevant variables and create a method for an investigation to address
your research questions?
Identify independent, dependent and controlled variables?
Undertake independent investigations safely and with competence?
Can you create and/or use a correct data table? This includes correct labels, units and consistent use of
decimal places.
Can you gather and record data in a safe and accurate manner?
Can you analyse your data and communicate it in an appropriate manner?
Uses critical thinking skills to evaluate information and drawing appropriate conclusions?
Can you evaluate weaknesses and errors in your method based on experimental evidence?
Can you?/Do you?
Understanding
Use internationally recognized symbols to write the formulae for common compounds?
Write and balance equations for a range of common chemical reactions?
Describe the meanings of the words solute, solvent and solution?
Define concentration of solutions using molarity (moles of solvent/dm3 of solution)?
Problem Solving
Calculate stoichiometric reacting mass and volumes of solutions using solutions concentrations
expressed in g/dm3 and mol/dm3? Can you…
● Express solution concentration in mol/dm3?
● Express solution concentration in g/dm3?
●
●
●
●
Calculate volumes of solutions given mass or moles?
Calculate moles of solutions given volume and/or mass?
Calculate moles or concentration of substance in a titration?
Determine the volumes of solutions needed to create dilutions using the dilution equation,
CCVC = CDVD
Solve problems involving limiting reagents?
Experimental Skills
105
Perform a titration (either to standardize a solution or determine the concentration of an unknown
solution)?
Calculate percent yield?
Can You/Do You…? (These objectives are for both atomic and electronic structure)
Understanding
●
●
●
●
●
●
Explaining
●
●
●
●
●
●
Language
●
Do you know that covalent bonding occurs between nonmetal atoms?
Do you know that covalent bonding involves sharing valence electrons so that each
atom obtains a noble gas configuration?
Do you know the names and formulas for the common covalent compounds water
(H2O), ammonia (NH3), carbon dioxide (CO2) and carbon monoxide (CO)?
Do you know the formulas for all of the diatomic elements—hydrogen (H2), oxygen
(O2), nitrogen (N2), fluorine (F2), chlorine (Cl2), bromine (Br2) and iodine (I2)—and
the most common forms of sulphur (S8) and phosphorus (P4)?
Do you know the different intermolecular forces: London dispersion forces (LDF),
dipole-dipole interaction (DP-DP) and hydrogen bonding (H-bond) and their relative
strengths?
Do you know that electronegativity (χ) is a measure of the attraction of an element
for the shared electrons in a covalent bond and that the higher the electronegativity
the stronger the attraction for the shared electrons is?
Can you use a labeled diagram to describe how covalent bonding occurs? This
should include both Bohr diagrams and Lewis diagrams.
Can you draw Lewis diagrams for simple covalent compounds? (Extension = Draw
Lewis diagrams for the common polyatomic ions.)
Can you write names and formulas for simple (two element) covalent compounds?
Can you use electronegativity difference (Δχ) to determine whether a bond is polar
or not (Δχ > 0.4)?
Can you determine the dominant intermolecular force for an atom or molecule
based on its structure?
Can you relate the strength of interparticle and/or intermolecular forces to melting
and boiling points?
Use a labeled diagram and a short paragraph to describe how covalent bonding
occurs. At a minimum you should use the following words/phrases: nonmetal,
share, valence electron, noble gas configuration, electrostatic attraction.
106
Lesson 26: Unit 3 Lesson 1 - Lewis Structures and Formal Charges
[Top of Doc]
Learning Goals
Problem Solving
●
●
●
Explaining
●
●
●
Can you draw Lewis diagrams for covalent compounds and polyatomic ions?
Can you assign formal charges to each element in a molecule or polyatomic
ion?
Can you determine the best orientation of electrons in a molecule or polyatomic
ion using the octet rule and formal charges?
Can you use a labeled diagram to describe how covalent bonding occurs? This
should include both Bohr diagrams and/or Lewis diagrams.
Can you explain why one Lewis structure is more stable than another using the
octet rule and formal charges?
Can you use a labeled diagram and write a short paragraph to describe how
covalent bonding occurs. At a minimum you should use the following
words/phrases: nonmetal, share, valence electron, noble gas configuration,
electrostatic attraction?
Drawing Lewis Diagrams
Fool Proof Lewis Diagrams:
1. Count the number of total valence electrons in your target molecule. Remember that if your molecule
is a polyatomic anion you must add electrons to your total (because anions have “extra” electrons)
and if your molecule is a polyatomic cation you must subtract electrons from your total (because
cations have “fewer” electrons).
2. Connect the atoms in your target molecule together with single bonds. Remember that if there is only
one of a particular element that element is likely to be the central atom.
3. Add electron pairs to your structure in a methodical way until you reach the correct total number.
4. If all of the atoms in your structure have 8 valence electrons (except for H which should have 2) you
are finished, if not move electron pairs from atoms that are “rich” in electrons to form additional
bonds with atoms that are “poor “ in electrons until every atom (except H) has 8 valence electrons
around it.
5. As a final check:
a. Count the valence electrons in your structure and ensure that it is the same as the total
number you counted in step 1.
b. Make sure that every atom (except H) has 8 valence electrons around it (a “noble gas
configuration”).
6. Double check that the sum of all formal charges correspond to the overall charge of the molecule or
ion.
----------------------------------------------------------------------------------------------------------------------------------------------
107
108
109
110
Assigning Formal Charges
Formal charges are charges we assign to each atom in a Lewis structure. This is not to be confused
with the net charge of an ion.
For example, the nitrate ion, NO3− has a net charge of −1. The N atom has a formal charge of +1
and each oxygen atom that is singly-bonded to N has a formal charge of −1.
Why Do We Need Formal Charges?
Formal charges can help us to determine if a Lewis Structure is correctly drawn. Often it is useful to
know exactly which part of the molecule that is charged. If you look at the
formula of the nitrate ion, NO3−, you only know that the net charge is −1 but you don’t know where it
is. The formula NO3- is not implying that the −1 is on the O.
Furthermore, the formal charges give us a means to determine whether a particular structure is
reasonable; or, if you have several resonance structures, the formal charges will help you determine
which resonance structure is preferred over the others. For example, you would not like to have a
111
structure with a positive formal charge on a highly electronegative element (such as oxygen). You
would also avoid writing a Lewis structure with a lot of formal charges (although in some cases it
cannot be avoided). Often it is preferred to violate the Octet Rule in order to achieve a zero formal
charge on an atom.
1. Examine each atom in the Lewis structure, one at a time.
2. Count both electrons in a lone pair (nonbonding electrons) and one electron per bond.
3. Compare this number with the valency/Group # (from the periodic table), which tells you how
many the atom is supposed to have.
4. The Group # tells you how many electrons a neutral atom would have. So, if you have one
more electron than the Group # indicates, the atom has a formal charge of −1. If you have two
more electrons, the charge is −2, and so forth.
5. If you have one less electron than the Group # indicates, the atom has a charge of +1. If you
have two less electrons, the charge is +2, and so forth.
Example: The nitrogen below has a formal charge of +1 because it is in possession of 4 electrons
using this formal charge technique but it has a valency of 5. This means it’s one short, so it is
assigned a formal charge of +1.
Determine the formal charges in each of the following.
2.
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3.
4. With reference to the following three structures (ignoring the arrows), rank them from most stable
to least stable based on formal charge rules. Hint: Assign formal charges first.
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Polyatomic Ion Structures and Formal Charges
Assign Formal Charges:
Assign Formal Charges:
Assign Formal Charges:
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Assign Formal Charges:
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Lesson 27: Unit 3 Lesson 2 - VSEPR Theory (1-4 Domains) [Top of Doc]
Learning Goals
Problem
Solving
Understanding
Can you determine the best orientation of electrons in a molecule or polyatomic ion using the octet
rule and formal charges?
Can you count the correct number of electron domains around a central atom?
Can you determine the geometric shape of molecules or polyatomic ions using VSEPR Theory (up
to 4 electron domains)?
Can you determine the bond angles in molecules or polyatomic ions using VSEPR Theory (up to 4
electron domains)?
Can you interpret stereochemical formulas of molecules or polyatomic ions after the shape and
bond angles have been determined using VSEPR theory?
Do you understand that VSEPR theory is based on the number of electron domains around a
central atom and the repulsion forces between electrons?
VSEPR Theory Practice
1. Explain what must occur if there isn’t enough valence electrons to fill all the outer electron shells
in order to achieve full octets?
2) For the following molecules/polyatomic ions:
i) Draw the Lewis dot diagram. Assign Formal Charges.
iii) Circle the electron domains. State the number.
ii) Predict the shape and the bond angle
a) HCN
b) PF3
c) COCl2
d) H2S
e) CO2
f) CO32–
g) SO32–
h) C2H4
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i) H2CO
j) NO3-
k) PH3
l) NO2-
m) POCl3
n) ammonia
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Lesson 28: Unit 3 Lesson 3
Review of IMF and Determination of Polarity
[Top of Doc]
Learning Goals
Problem
Solving
Understanding
Identifying Intermolecular Forces (REVIEW from Grade 9)
a. Put an “X” into the box for every intermolecular force that element or compound has.
i. LDF = London dispersion force
ii. DP-DP = Dipole-dipole interaction
iii. H-Bond = Hydrogen bond
b. Identify the dominant intermolecular force for that element or compound.
c. Briefly explain why you identified that particular dominant force. You may need to refer to the
electronegativity table at the end of this packet.
See the example below in the first row.
Element or
LDF DP- H-Bo Dominant
Brief explanation of dominant IMF
Compound
DP
nd
IMF
H is bonded directly to F (a highly electronegative
HF
X
X
X
H-Bond
atom) so it has hydrogen bonding
PCl3
C3H8
CH3CH2NH2
CH3OCH3
2. Which compound has the highest melting point?
a. CH3CH2CH3
b. CH3CH2OH
c. CH3OCH3
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d. CH3CHO
3. Briefly explain your answer to question #2.
4. What is the correct order if the compounds are arranged in order of increasing boiling point?
a. CH4 < CH3Cl < SiH4 < CH3OH
b. CH3OH < CH4 < CH3Cl < SiH4
c. CH3OH < CH3Cl < SiH4 < CH4
d. CH4 < SiH4 < CH3Cl < CH3OH
5. Briefly explain your answer to question #3.
6. The graph below shows the boiling points of the hydrides of group 5. Use it to answer the questions
below.
a. What happens to the boiling points as you go down the group, starting at phosphorus?
b. Briefly explain why this occurs from PH3 to SbH3.
c. Briefly explain why NH3 has such an unusually high boiling point.
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7. Ethene, C2H4, and hydrazine, N2H4, are hydrides of adjacent elements on the periodic table.
a. Draw Lewis structures for both compounds below.
Ethene
Hydrazine
b. The polarity of a molecule can be defined using the concept of electronegativity.
i. What is electronegativity?
ii. Compare the electronegativity differences of the C-H bond and the N-H bond.
Ethene
Hydrazine
iii. Identify each compound as being polar or non-polar and briefly explain your
identification.
Ethene
Hydrazine
iv. Ethene has a boiling point of -107oC while hydrazine has a boiling point of 114oC.
Briefly explain this difference in terms of their intermolecular forces.
Determining if Molecules are Polar (New)
1) Without the use of a text or your notes, create a flowchart below to summarize the steps involved
in determining whether a molecule will be polar or non-polar. Include a few molecules along the way
to make sure to reinforce each step of the process.
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2) With the aid of a diagram, explain how a molecule can have polar bonds but still be considered
non-polar in nature.
……………………………………………………………………………………………………
……………………………………………………………………………………………………...
3) Draw the correct dipole arrow for the bond between hydrogen and nitrogen. What does this tell
you about their electronegativities?
……………………………………………………………………………………………………..
……………………………………………………………………………………………………...
4) Explain why Bromobenzene and benzene have different polarities. Benzene is a cyclic compound
with the molecular formula C6H6.
………………………………………………………………………………………………………
……………………………………………………………………………………………………...
5) Explain why water is considered polar despite having the same atoms bonded to oxygen.
………………………………………………………………………………………………………
……………………………………………………………………………………………………...
6) Explain why tetrachloromethane is considered non-polar even though it has polar bonds.
………………………………………………………………………………………………………
……………………………………………………………………………………………………...
7) Deduce the shape and bond angles of:
(a) boron trichloride, BCl3
………………………………………………………………………………………………………
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………………………………………………………………………………………………………
i) Is boron trichloride polar or non-polar? Explain.
………………………………………………………………………………………………………
………………………………………………………………………………………………………
(b) phosphoryl chloride, POCl3
………………………………………………………………………………………………………
………………………………………………………………………………………………………
(c) phosphine, PH3
………………………………………………………………………………………………………
……………………………………………………………………………………………………...
(d) hydrogen cyanide, HCN
………………………………………………………………………………………………………
……………………………………………………………………………………………………...
9) Explain why hydrogen fluoride, HF, and water, H2O, are very polar molecules but
tetrafluoromethane, CF4, and carbon dioxide, CO2 are non- polar.
………………………………………………………………………………………………………
………………………………………………………………………………………………………
………………………………………………………………………………………………………
………………………………………………………………………………………………………
10) Fill in the table below:
Compound
Lewis D.D.
Bond Type
NPC/PC/I
Shape
Symmetrical
Y/N
Molecular
Polarity
P or NP
CH4
CO2
NH3
NCl3
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CF4
CCl4
H2CO
NaCl
H2O
Lesson 29: Unit 3 Lesson 4 - Chromatography (Polarity and Solubility)
[Top of Doc]
Learning Goals
Problem
Solving
Understanding
Solubility Demonstration
Previewing question:
Are the following substances polar or nonpolar?
Water
Propanone
Hexane
Propanol
Optional Watch: 38 Solubility Demonstration (9:21)
Record whether or not each of the following mixtures were homogeneous or heterogeneous.
Water + Propanone
Water + Hexane
Water + Propanol
In the space below sketch a diagram of each mixture.
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Water + Propanone
Water + Hexane
Water + Propanol
Fill in the blanks with the words “soluble”, “insoluble” or ”slightly soluble”
Propanone is _________________________________ in water.
Hexane is ___________________________________ in water.
Propanol is __________________________________ in water.
What is the general solubility rule Mr Smith mentioned in the video.
A separatory funnel can be used to separate heterogeneous mixtures of liquids. In the water + hexane
mixture what substance formed the top layer and which substance formed the bottom layer? Explain this
observation.
What 2 physical properties does a separatory funnel use?
Consider the bromine molecule shown in the box below.
Would you expect bromine to be more soluble in water or hexane? Explain.
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Consider the ethanoic acid molecule shown in the box below.
Would you expect ethanoic acid to be more soluble in water or hexane? Explain.
Thin-Layer Chromatography Worksheet
1. Chromatography usually involves two different phases. What names are given to these two phases?
2. Shown below is a chromatogram of three unknown amino acids, A, B, and C, and a reference table of
𝑅f values of various amino acids. What are the three amino acids?
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Amino Acid
Lysine
Glycine
Alanine
Proline
Leucine
𝑅 value
0.13
0.26
0.38
0.43
0.73
3. A student is running a chromatography experiment. The diagram below shows their setup at the start
of the experiment. What mistake has the student made?
a) A lid has been placed over the beaker
b) There is no thermometer to measure the temperature.
c) The solvent level is above the origin.
d) The spots are not below the origin.
4. Which of the following is the main function of TLC?
a) Separating a mixture of a liquid and gas
b) Separating a mixture of gases
c) Separating a mixture of a solid and liquid
d) Separating a mixture of liquids
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5. The given chromatogram shows that substance B has traveled further up the chromatography paper
than A. Which properties of a substance affect how far it will travel up the chromatography paper?
a) Its density compared to the density of the paper (stationary phase).
b) Its solubility in the solvent (mobile phase), and its attraction to the paper
(stationary phase)
c) Its boiling point and volatility.
d) Its molar mass.
6. Consider Vitamins A and C whose structures are shown below:
These two vitamins can be separated, using TLC, by different solvents. For example, one could use either
ethanol, CH3CH3OH, or octane (C8H18).
a) In one of these solvents vitamin A has an Rf value of 0.15 and vitamin C has an Rf value of 0.96.
Which solvent was used? Explain your reasoning.
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b) On the diagram below show where you would find vitamin A? Include a calculation to support your
answer.
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[Top of Doc]
Lesson 30: Unit 3 Lesson 5
Formative Lab 1 - Polarity and Chromatography
6
Separating Mixtures: Thin-Layer Chromatography Lab
Theory:
Chromatography is probably the most useful method of separating organic compounds for identification and
purification. There are many different types of chromatography but most work on the concept of absorbance.
The two important parts of chromatography are the adsorbent (stationary phase) and the eluent (mobile
phase or solvent). A good adsorbent is usually a solid material that will attract and absorb the materials to
separated. Paper, silica gel or alumina are all very good absorbents. The eluent is the solvent which carries
the a materials to be separated through the adsorbent.
Chromatography works on the concept that the compounds to be separated are slightly soluble in the eluent
and will spend some of the time in the eluent (or solvent) and some of the time on the absorbent. When the
compounds to be separated spend different amounts of time being absorbed, then they can be separated
from one another. The polarity of the molecules to be separated and the polarity of the eluent are very
important. Changing the polarity of the eluent will only slightly change the solubility of the molecules but will
greatly change their willingness or affinity to absorb onto the absorbent. This affinity for the eluent versus the
absorbent is what separates the molecules.
A simple example of chromatography is the separation of the various dyes in ink using paper
chromatography. In this example, paper is the absorbent and water is the eluent. The various dyes spend
differing amounts of time absorbing on the paper due to their solubility in water and will begin to separate as
the water carries them further from the starting point.
To separate complex organic molecules, thin layer chromatography (TLC) is frequently used. In TLC, the
absorbent is usually silica gel (SiO2) or alumina (Al2O3) and the eluent is an organic solvent. The polarity of
the eluent is very important in TLC since a small change in polarity can dramatically increase or decrease the
solubility of some organic molecules. Many times, a mixture of a nonpolar solvent (hexane) and a polar
solvent (acetone) are used to achieve an optimum polarity.
Mixture components that are separated through TLC are often identified using Rf values. Rf values are
calculated by dividing the distance that the component travels from its initial location (A in the diagram below)
by the distance that the solvent has travelled (the solvent front-S in the diagram below). These values can be
used to compare mixture components and also to identify them.
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Safety:
Acetone is flammable and a dangerous fire risk, toxic by ingestion and inhalation. Hexanes are flammable
and a dangerous fire risk, may be irritating to respiratory tract. Always wear chemical splash goggles.
Objectives:
● To separate a mixture using thin-layer chromatography
● To compare plant pigments by Rf values
Equipment:
● 250 mL beaker
● 10 mL measuring cylinder
● test tube x 2
●
●
●
plastic TLC plate
capillary tube, open-end
test tube rack x 2
●
●
●
Pasteur pipet x 2
plastic wrap
Parafilm
Materials:
● spinach
● anhydrous magnesium sulphate – MgSO4 (s)
● fine sand
● acetone
● acetone/hexane (1:4) developing solution
Procedure A: Sample Preparation - ALREADY COMPLETED BY LAB TECHNICIAN
The following procedure could be carried out for both the spinach and the local plant leaf. Each lab partner
should prepare a sample.
1. Weigh out 0.5 g of the sample. Combine this with 0.5 g of anhydrous magnesium sulphate and 1.0 g
of sand.
2. Using a mortar and pestle, grind the mixture until it become a fine, light green powder.
3. Transfer this solid into a test tube and 2 mL of acetone. You should use one of the Pasteur pipets to
transfer acetone to the 10 mL graduated cylinder for measurement.
4. Stir the solution with the stirring rod for 2 minutes and allow it to sit for 10 minutes. The solid should
settle to the bottom, leaving a green liquid layer on top. WHILE YOU ARE WAITING begin procedure
B.
5. Transfer this green layer to another test tube using a pipet and seal the test tube containing the
extract with Parafilm.
Alternate Procedure for Sample Preparation:
Your teacher may choose to provide you with pre-prepared samples in order to save time. These samples are
prepared with the same acetone solvent and magnesium sulphate but a blender has been used to speed up
preparation.
Procedure B: TLC of the Extract - STUDENTS START HERE
39 Thin Layer Chromatography (6:30)
1. Fill the developing tank (small beaker) with the developing solution to a depth of approximately 0.5 cm
and cover with plastic film. This will ensure that the atmosphere is saturated with the developing
solution, a 1:4 solution of acetone:hexane.
2. Prepare the TLC plate by using a ruler to draw a light pencil line across the bottom of the plate. The
line MUST be higher than the depth of the developing solution in the beaker.
3. Spot the TLC plate by briefly and gently touching the tip of the capillary tube containing the green
acetone solution to the TLC plate on the pencil line, gently blowing on the spot to evaporate the
acetone and repeating the process several times until you have built up a dark green spot.
4. Place the TLC plate in the development tank making sure that the spot is above the solvent level and
quickly cover the tank with plastic film.
5. Observe the solvent as it travels up the TLC plate and begins to separate the components from the
spinach solution. When the solvent almost reaches the top of the plate, remove the plate from the
tank and allow it to dry.
6. Observe and record any coloured spots on the plate. A yellow spot near the top and several greenish
spots nearer the bottom should be present.
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Procedure B: Analysis of the TLC
1. Carefully measure the distance from the pencil line to the solvent front. This is S.
2. For each spot on the TLC measure the distance from the pencil line to the center of that spot. This is
A for each spot.
3. Calculate the Rf value for each spot (Rf = A/S).
4. Make a careful drawing of your TLC plate and label it with all of your measured and calculated values.
Data Drawing: Show your measurements and calculated values
Interpretation
If done correctly, the TLC will normally reveal five to six different pigments. The pigments can be identified by
their colours and relative positions on the chromatogram. The major pigments are (in order from the initial
spot):
1. chlorophyll b (olive green)
2. lutein (yellow)
3. chlorophyll a (blue green)
4. violaxanthin (yellow)
5. pheophytin b (grey)
6. carotene (yellow-orange), near the solvent front
Other pigments may be visible but these are more difficult to identify. Sometimes neoxanthin is also present
but is usually very closer to the chlorophyll b and not visible.
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Lesson 32: Unit 3 Lesson 7
Introduction to Redox Processes
[Top of Doc]
Lesson 33: Unit 3 Lesson 8
Balancing Redox Equations
[Top of Doc]
Lesson 34: Unit 3 Lesson 9
Voltaic Cells
[Top of Doc]
Lesson 35: Unit 3 Lesson 10
Formative Quiz 2/Review
[Top of Doc]
Lesson 37: Cumulative Design Lab - Day 1
[Top of Doc]
Lesson 38: Cumulative Design Lab - Day 2
[Top of Doc]
Lesson 39: Exam Corrections and Lab Peer-Assessment [Top of Doc]
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