Chemistry 20 Plan
Course Outline
Chemistry 20 is the second course towards completion of the chemistry program from
Alberta Education. This introductory chemistry course builds on basic content started in
Science 10 and introduces students to the topics required to describe chemicals and
chemical systems. Successful completion of Chemistry 20 serves as a prerequisite for
Chemistry 30. The chemistry program is designed for students planning a postsecondary education. It is also recommended that students taking chemistry have a
strong math background.
Prerequisite: Science 10 (recommended mark of 65%)
Textbook:
Nelson Chemistry (2007)
Evaluation & Time Distribution:
Science 10 Review
3%
2 weeks
Chemical Bonding
6%
3 weeks
Gases
5%
2 weeks
Solutions
9%
4 weeks
Cumulative Exam
10%
Stoichiometry
12%
Labs, Quizzes, Assignments
25%
Final Exam
30%
TOTAL
100%
4 weeks
18 (16 ps3) weeks
Left three days before Christmas as flex time in the course (extra time to work with)
Units:
A. The Diversity of Matter and Chemical Bonding
B. Forms of Matter: Gases
C. Matter as Solutions, Acids and Bases
D. Quantitative Relationships in Chemical Changes
Meeting the ICT outcomes
Overview: Meeting the ICT outcomes
Computers and technology are integral to our modern way of life. The ICT outcomes promote
computer and technological literacy. Science utilizes computers for data analysis as well as data
presentation. Students will use technology to record experiments and analyze data in a useful fashion
effectively integrating computer technology into the curriculum.
Requisite Science 10 Knowledge outcomes applicable to Chemistry 20
Science 10: Unit A : Energy in Matter and Chemical Change
Learning Outcomes:
Students will:
1. Describe the basic particles that make up the underlying structure of matter, and investigate
related technologies
• identify historical examples of how humans worked with chemical substances to meet their basic needs
(e.g., how pre-contact First Nations communities used biotic and abiotic materials to meet their needs)
• outline the role of evidence in the development of the atomic model consisting of protons and neutrons
(nucleons) and electrons; i.e., Dalton, Thomson, Rutherford, Bohr
• identify examples of chemistry-based careers in the community (e.g., chemical engineering,
cosmetology, food processing)
2. Explain, using the periodic table, how elements combine to form compounds, and follow IUPAC
guidelines for naming ionic compounds and simple molecular compounds
• illustrate an awareness of WHMIS guidelines, and demonstrate safe practices in the handling, storage
and disposal of chemicals in the laboratory and at home
• explain the importance of and need for the IUPAC system of naming compounds, in terms of the work
that scientists do and the need to communicate clearly and precisely
• explain, using the periodic table, how and why elements combine to form compounds in specific ratios
• predict formulas and write names for ionic and molecular compounds and common acids (e.g., sulfuric,
hydrochloric, nitric, ethanoic), using a periodic table, a table of ions and IUPAC rules
• classify ionic and molecular compounds, acids and bases on the basis of their properties; i.e.,
conductivity, pH, solubility, state
• predict whether an ionic compound is relatively soluble in water, using a solubility chart
• relate the molecular structure of simple substances to their properties (e.g., describe how the properties
of water are due to the polar nature of water molecules, and relate this property to the transfer of energy
in physical and living systems)
• outline the issues related to personal and societal use of potentially toxic or hazardous compounds (e.g.,
health hazards due to excessive consumption of alcohol and nicotine; exposure to toxic substances;
environmental concerns related to the handling, storage and disposal of heavy metals, strong acids,
flammable gases, volatile liquids)
3. Identify and classify chemical changes, and write word and balanced chemical equations for
significant chemical reactions, as applications of Lavoisier’s law of conservation of mass
•provide examples of household, commercial and industrial processes that use chemical reactions to
produce useful substances and energy (e.g., baking powder in baking, combustion of fuels, electrolysis of
water into H2(g) and O2(g))
• identify chemical reactions that are significant in societies (e.g., reactions that maintain living systems,
such as photosynthesis and respiration; reactions that have an impact on the environment, such as
combustion reactions and decomposition of waste materials)
• describe the evidence for chemical changes; i.e., energy change, formation of a gas or precipitate, colour
or odour change, change in temperature
• differentiate between endothermic and exothermic chemical reactions (e.g., combustion of gasoline and
other natural and synthetic fuels, photosynthesis)
• classify and identify categories of chemical reactions; i.e., formation (synthesis), decomposition,
hydrocarbon combustion, single replacement, double replacement
• translate word equations to balanced chemical equations and vice versa for chemical reactions that occur
in living and nonliving systems
• predict the products of formation (synthesis) and decomposition, single and double replacement, and
hydrocarbon combustion chemical reactions, when given the reactants
• define the mole as the amount of an element containing 6.02 × 1023 atoms (Avogadro’s number) and
apply the concept to calculate quantities of substances made of other chemical species (e.g., determine the
quantity of water that contains 6.02 × 1023 molecules of H2O)
• interpret balanced chemical equations in terms of moles of chemical species, and relate the mole concept
to the law of conservation of mass
Crucial Questions from Science 10 Knowledge outcomes
Student understanding regarding the following 16 questions will be covered in the first two weeks of
class as per the course outline for chemistry 20 at Winston Churchill. These questions were chosen
based upon their concept inclusion within the science 10 chemistry unit and their importance with
regard to building student understanding of chemistry going forward. Two to three questions will be
addressed per class in the order stated below.
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
What are the WHMIS symbols and what do they mean?
What is all matter composed of?
What subatomic particles make up the atom?
Which subatomic particles are involved in most/all chemical reactions?
How is the Periodic Table organized?
What happens in ion formation, with regard to valence electrons?
How scientists name ionic and molecular compounds?
What is a chemical reaction?
Classify and balance the types of chemical reactions?
Exothermic and Endothermic?
11.
12.
13.
14.
15.
16.
What are acids and bases?
What does solubility mean?
How do we use a solubility table?
What is molar mass and how do we find it?
What does mole mean, or Avagadro’s number?
How does mole have to do with conservation of mass?
Tentative First Two Week Schedule
Monday – Sep 1
Holiday
Sep 8
ROUND 2 TEXT
BOOKS
Quiz:
(Periodic table
structure,
electron
configuration,
level diagram)
Naming
Compounds
Chemical Reaction
Characteristics
Types of Chemical
Reactions,
Prediction of
reactions
exothermic and
Endothermic
TASKS:
Tuesday -2 classes
Sep 2
Meet and greet
the students
Explain Chemistry
20 course
according to
outline.
Expectations and
Rules
Wednesday
Sep 3 – day 2
GET TEXTBOOKS
WHMIS Symbols
Atoms, atomic
Structure,
(Valence shell
electrons)
Solids Liquids and
gases
Periodic Table
Structure. (groups
and SPDF)
Handouts:
Periodic Table
Assignment 1
Thursday double
Sep 4
Periodic table
structure – SPDF
block structure,
counting electrons
Electron
configurations
Electron
shell/level
diagrams
TASKS:
Mark Assignment
1
Handouts:
Assignment 2
Friday
Sep 5
Quiz:
(WHMIS Symbols
quiz + Periodic
Table Structure)
Ion Formation
Effect upon
energy level
diagrams and
electron
configuration.
TASKS:
Mark Assig 2
Handout:
Assignment 3
Sep 9
Quiz:
(Ion Formation,
Energy level
diagrams of ions
Electron
configuration of
Ions)
Acid and Base
characteristics
Balancing
Chemical
Reactions
Solubility Tables
TASKS:
Mark assign 4
Handouts:
Assign 5
Sep 10
Quiz:
(Characteristics of
Chemical
Reactions,
classification,
product
prediction, endo
and exothermic)
Avagadro’s
number moles,
molar mass,
Conservation of
Mass.
TASKS:
Mark Assign 5
Handouts:
Assign 6
Sep 11
Quiz (acids bases,
reaction
balancing,
solubility)
Task:
Mark assign 6
Review of
Problem Areas
Review package
Sep 12
Review of Science
10 Chemistry Test
Mark Assign 3
Handout: Assign 4
Materials Preparation for first two weeks –
1. Periodic Table, 6 assignments, 5 quizzes, Science 10 review test
On to Chem 20 material
Chemistry 20 learning outcomes from the program of studies
Unit A: The Diversity of Matter and Chemical Bonding
Themes: Diversity and Matter
Overview: Concepts, models and theories are often used in interpreting and explaining observations
and in predicting future observations. The major focus of this unit is to relate theories about bonding to
the properties of matter and to develop explanations and descriptions of structure and bonding through
scientific models. Students learn about the diversity of matter through the investigation of ionic
compounds and molecular substances. This unit should be complete by:
Key Concepts: The following concepts are developed in this unit and may also be addressed in
other units or in other courses. The intended level and scope of treatment is defined by the
outcomes.
• chemical bond
• ionic bond
• covalent bond
• electronegativity
• polarity
• valence electron
• intramolecular and intermolecular forces
• hydrogen bond
• electron dot diagrams
• Lewis structures
• valence-shell electron-pair repulsion (VSEPR) theory
Focusing Questions
Learning Outcomes
GLO 1: Students will describe the role of modelling, evidence and theory in explaining and
understanding the structure, chemical bonding and properties of ionic compounds.
20–A1.1k recall principles for assigning names to ionic compounds
20–A1.2k explain why formulas for ionic compounds refer to the simplest whole-number ratio of ions
that result in a net charge of zero
20–A1.3k define valence electron, electronegativity, ionic bond and intramolecular force
20–A1.4k use the periodic table and electron dot diagrams to support and explain ionic bonding theory
20–A1.5k explain how an ionic bond results from the simultaneous attraction of oppositely charged ions
20–A1.6k explain that ionic compounds form lattices and that these structures relate to the compounds’
properties; e.g., melting point, solubility, reactivity.
GLO 2: Students will describe the role of modelling, evidence and theory in explaining and
understanding the structure, chemical bonding and properties of molecular substances.
20–A2.1k recall principles for assigning names to molecular substances
20–A2.2k explain why formulas for molecular substances refer to the number of atoms of each
constituent element
20–A2.3k relate electron pairing to multiple and covalent bonds
20–A2.4k draw electron dot diagrams of atoms and molecules, writing structural formulas for molecular
substances and using Lewis structures to predict bonding in simple molecules
20–A2.5k apply VSEPR theory to predict molecular shapes for linear, angular (V-shaped, bent),
tetrahedral, trigonal pyramidal and trigonal planar molecules
20–A2.6k illustrate, by drawing or by building models, the structure of simple molecular substances
20–A2.7k explain intermolecular forces, London (dispersion) forces, dipole-dipole forces and hydrogen
bonding
20–A2.8k relate properties of substances (e.g., melting and boiling points, enthalpies of fusion and
vaporization) to the predicted intermolecular bonding in the substances
20–A2.9k determine the polarity of a molecule based on simple structural shapes and unequal charge
distribution
20–A2.10k describe bonding as a continuum ranging from complete electron transfer to equal sharing of
electrons.
Unit B: Forms of Matter: Gases
Themes: Matter, Change and Energy
Overview: Students expand their knowledge of the nature of matter through the investigation of the
properties and behaviour of gases.
Key Concepts: The following concepts are developed in this unit and may also be addressed in
other units or in other courses. The intended level and scope of treatment is defined by the outcomes.
• Celsius and Kelvin temperature scales
• Boyle’s law
• ideal gas law
• standard temperature and pressure (STP)
• standard ambient temperature and pressure (SATP)
• absolute zero
• real and ideal gases
• law of combining volumes
• Charles’s law
FOFOCUSING QUESTIONS



How do familiar observations of gases relate to specific scientific models describing the
behaviour of gases?
What is the relationship among the pressure, temperature, volume and amount of a gas?
How is the behaviour of gases used in various technologies?
GLO 1: Students will explain molecular behaviour, using models of the gaseous state of matter.
20–B1.1k describe and compare the behaviour of real and ideal gases in terms of kinetic molecular theory
20–B1.2k convert between the Celsius and Kelvin temperature scales
20–B1.3k explain the law of combining volumes
20–B1.4k illustrate how Boyle’s and Charles’s laws, individually and combined, are related to the ideal
gas law (PV = nRT)
• express pressure in a variety of ways, including units of kilopascals, atmospheres and millimetres of
mercury
• perform calculations, based on the gas laws, under STP, SATP and other defined conditions.
Unit C: Matter as Solutions, Acids and Bases
Themes: Matter, Diversity, Systems and Change
Overview: Students gain insight into the nature of matter through an investigation of change in the
context of solutions, acids and bases.
Key Concepts: The following concepts are developed in this unit and may also be addressed in
other units or in other courses. The intended level and scope of treatment is defined by the
outcomes.
• homogeneous mixtures
• solubility
• electrolyte/nonelectrolyte
• concentration
• dilution
• strong acids and bases
• weak acids and bases
• monoprotic/polyprotic acid
• monoprotic/polyprotic base
• Arrhenius (modified) theory of acids and bases
• indicators
• hydronium ion/pH
• hydroxide ion/pOH
• neutralization
GLO 1: Students will investigate solutions, describing their physical and chemical properties.
20–C1.1k recall the categories of pure substances and mixtures and explain the nature of homogeneous
mixtures
20–C1.2k provide examples from living and nonliving systems that illustrate how dissolving substances
in water is often a prerequisite for chemical change
20–C1.3k explain dissolving as an endothermic or exothermic process with respect to the breaking and
forming of bonds
20–C1.4k differentiate between electrolytes and nonelectrolytes
20–C1.5k express concentration in various ways; i.e., moles per litre of solution, percent by mass and
parts per million
20–C1.6k calculate, from empirical data, the concentration of solutions in moles per litre of solution and
determine mass or volume from such concentrations
20–C1.7k calculate the concentrations and/or volumes of diluted solutions and the quantities of a solution
and water to use when diluting
20–C1.8k use data and ionization/dissociation equations to calculate the concentration of ions in a
solution
20–C1.9k define solubility and identify related factors; i.e., temperature, pressure and miscibility
20–C1.10k explain a saturated solution in terms of equilibrium; i.e., equal rates of dissolving and
crystallization
20–C1.11k describe the procedures and calculations required for preparing and diluting solutions.
GLO 2: Students will describe acidic and basic solutions qualitatively and quantitatively.
20–C2.1k recall International Union of Pure and Applied Chemistry (IUPAC) nomenclature of acids and
bases
20–C2.2k recall the empirical definitions of acidic, basic and neutral solutions determined by using
indicators, pH and electrical conductivity
20–C2.3k calculate H3O+(aq) and OH–(aq) concentrations and the pH and pOH of acidic and basic
solutions based on logarithmic expressions; i.e., pH = –log[H3O+] and pOH = –log[OH–]
20–C2.4k use appropriate Système international (SI) units to communicate the concentration of solutions
and express pH and concentration answers to the correct number of significant digits; i.e., use the number
of decimal places in the pH to determine the number of significant digits of the concentration
20–C2.5k compare magnitude changes in pH and pOH with changes in concentration for acids and bases
20–C2.6k explain how the use of indicators, pH paper or pH meters can be used to measure H3O+(aq)
20–C2.7k define Arrhenius (modified) acids as substances that produce H3O+(aq) in aqueous solutions
and recognize that the definition is limited
20–C2.8k define Arrhenius (modified) bases as substances that produce OH–(aq) in aqueous solutions and
recognize that the definition is limited
20–C2.9k define neutralization as a reaction between hydronium and hydroxide ions
20–C2.10k differentiate, qualitatively, between strong and weak acids and between strong and weak bases
on the basis of ionization and dissociation; i.e., pH, reaction rate and electrical conductivity
20–C2.11k identify monoprotic and polyprotic acids and bases and compare their
ionization/dissociation.
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
Recall differences between Pure Substance, Mixture and Homogeneous Mixtures (solutions)
Describe how dissolving substances in water is necessary for chemical reactions to take place
Describe the difference between Solvent and Solute
Describe dissolving as an exothermic or endothermic process, due to breaking and forming of
bonds
Differentiate between Electrolytes and Non-electrolytes
Express concentration Molarity, % mass, and parts per million (ppm)
Calculate the concentration of solution in moles per litre
Calculate the concentration and volume of solution dilutions (adding more solvent or what
concentration of original stock solution was)
Use dissociation equations to calculate the concentration of Ions in solution
Identify related solubility factors – Temperature, pressure, Miscibility
Explain that equilibrium of saturated solution is an equal amount dissolving and crystallizing
12. Describe the procedure for preparing solutions of known concentration by dilution or from pure
substance.
Second Topic - Acids and Bases
1. Recall naming of acids and bases
2. Recall definitions of acidic, basic and neutral solutions determined by using indicators (litmus
paper, pH and electrical conductivity
3. Calculate H3O+(aq) and OH–(aq) concentrations and the pH and pOH of acidic and basic solutions
4. Express pH and concentration to the right number of significant digits
5. Compare changes in pH and pOH with changes in concentration for acids and bases
6. Explain how indicators are used to test for acids
7. Define acids (Arrhenius definition) as substances that produce H3O+(aq) in aqueous solutions
8. Define Bases (Arrhenius definition) as substances that produce OH–(aq) in aqueous solutions
9. Define Neutralization as the reaction between hydronium and hydroxide
10. Differentiate, by calculation, between strong and weak acids and bases in relation to their
ionization and dissociation. pH, reaction rate and electrical conductivity.
11. Identify monoprotic and polyprotic acids and compare their dissociation and ionization.
Polyprotic can donate more than one H+(aq) to solution
Unit D: Quantitative Relationships in Chemical Changes
Themes: Matter and Change
Overview: Students focus on chemical change and the quantitative relationships contained in balanced
chemical equations. They are required to use stoichiometric principles and mathematical manipulation
to predict quantities of substances consumed or produced in chemical reaction systems.
Key Concepts: The following concepts are developed in this unit and may also be
addressed in other units or in other courses. The intended level and scope of treatment is
defined by the outcomes.
• chemical reaction equations
• net ionic equations
• spectator ions
• reaction stoichiometry
• precipitation
• limiting and excess reagents
• actual, theoretical and percent yield
• titration
• end point
• equivalence point
• titration curves for strong acids and bases
GLO 1: Students will explain how balanced chemical equations indicate the quantitative relationships
between reactants and products involved in chemical changes.
20–D1.1k predict the product(s) of a chemical reaction based upon the reaction type
20–D1.2k recall the balancing of chemical equations in terms of atoms, molecules and moles
20–D1.3k contrast quantitative and qualitative analysis
20–D1.4k write balanced ionic and net ionic equations, including identification of spectator ions, for
reactions taking place in aqueous solutions
20–D1.5k calculate the quantities of reactants and/or products involved in chemical reactions, using
gravimetric, solution or gas stoichiometry.
GLO 2: Students will use stoichiometry in quantitative analysis.
20–D2.1k explain chemical principles (i.e., conservation of mass in a chemical change), using
quantitative analysis
20–D2.2k identify limiting and excess reagents in chemical reactions
20–D2.3k define theoretical yields and actual yields
20–D2.4k explain the discrepancy between theoretical and actual yields
20–D2.5k draw and interpret titration curves, using data from titration experiments involving strong
monoprotic acids and strong monoprotic bases
20–D2.6k describe the function and choice of indicators in titrations
20–D2.7k identify equivalence points on strong monoprotic acid–strong monoprotic base titration
curves and differentiate between the indicator end point and the equivalence point.
Organization: I have presented the topics and the assignments I plan on introducing, along with the
weeks for each unit. I have not presented it as per day as I do not find this useful or beneficial as
students needs vary. Changes to assignments and topic material will change as needed to meet the
weakly knowledge goals. Small class starting quizzes that test cumulative knowledge are also not
presently shown as they tend to beginning of class exercises designed to start class on the right foot and
stimulate information recall in preparation for the final examination.
Schedule by Week
Week
Week 1
Sep 2 - 5
Unit
Review
Topics
Atomic Structure
Material layout
Atoms and ions
Ionic and molecular
compounds
Week 2
Sep 8 - 12
Review
Nomenclature
Compounds
Reactions
balancing
reactions
Week 3
Sep 15 – 19
Bonding
A1.1 – A1.5
A2.1 – A2.4
A2.10
Finish review
Introduction to
bonding
Key concepts
Intramolecular
bonding
Ionic and molecular
naming
The mole
Mathematical
conversions
Review Period
Introduction to
bonding and orbital
theory
Molecular modeling
Metallic Bonds
Electronegativity
Assignments
Worksheet bundle
+ Addison Wesley
black line masters
from science 10
Review Unit
Questions
Review Unit Test
Pg 78-84
Pg 84 – #4-10pg
Pg 85-90
Pg 89 #6, 90 #1 - 5
Worksheet
Bonding game
Ionic Compound
formulas
Molecular
compound
formulas
VSEPR theory and
molecular polarity
Intermolecular
bonding continued
Polarity
Lewis Structures
Week 4
Sep 22 - 26
Bonding
A2.5 - A2.9
Intermolecular
bonding (video on
intermolecular forces)
– dipole-dipole –
London forces
Hydrogen bonding
Week 5
Sep 29 – Oct 3
Bonding
A1.6
Ionic compounds:
Covalent Network
Review of bonding
Continuous lattice
Structures.
Ramifications of polar
bonds in covalent
network
Week 6
Oct 6 - 10
Gases
B1.1 – B1.3
Introduction to
gases, basic laws
Boyles Law
Charles Law
Combined gas law
Demonstration/work
Week 7
Oct 13 - 17
Gases
B1.4
Conditions/volume STP/SATP
calculations
Kinetic molecular
Gas volume
Mol / volume
Pg 96 1-4
Pg 104
#2,3,4,5,6,10
Quiz – Lewis ,
bond polarity,
compound
formulas
Worksheet
packet
Pushed into next
week
Page 110 –
Exercise #3A
Worksheet on
Intermolecular
forces
Pg 117 #1-5
Pg 118 # 12
Pg 119 - 129
Quiz:
Intermolecular
forces
Pg 122 # 1,7,10,11
Pg 128 #13,14,18
Network structure
Game
Unit Review
Questions Pg 138
#26,28,29,31,34,
38,39,41, 43,
47,48,49,50, 53 60
Bonding Test
Pg 146 – 157
Pg 150 2-4
Pg 152 6-10
Pg 156 14 -16, 18
Pg 161 #1-7
Worksheet
Quiz – gas laws
Pg 163 – 181
Pg 166 #5-7
Pg 171 # 6-14
relationships
Week 8
Oct 20 - 24
Solutions Acid
Base
Introduction to
solutions
Energy Changes
Aqueous solutions
Week 9
Oct 27 – Oct 31
Solutions Acid
Base
Concentration of
species in solution
Week 10
Nov 3 – 7
Solutions Acid
Base
Skill objectives
Week 11
Nov 10 – 14
(short week)
Solutions Acid
Base
Acids and bases
pH
Week 12
Nov 17 – 21
(short week)
Solutions Acid
Base
Hydronium
PH calculations
calculations
Ideal Gas Law
Lab – molar mass
Review
Molar volume
worksheet
Pg 176 9-12
Gas calculation
Worksheet2
Lab report
Review Pg 181 –
1-19
Chapter test
Solutions and mixtures Demo solutions
Dissociation/ionization Lab Exercise 5A
separation
Pg 195 #1-7
Electrolyte
Demo Conductivi
conductivity
Pg 197 -201
Litmus acid/base/neut Lab exercise 5B (
Review ionization and pg 202)
dissociation
Pg 202 #1-3, 8-10
Concentration
Pg 203 – 213
Ion concentration
Nelson worksheet
Ionization and
22
dissociation
Types 1-6
worksheet
Pg 212 #17-19
Pg 214 #12
Concentration
Practice
worksheet
Quiz
concentration
Solution preparation
Pg 215-219
Dilution
worksheet
Making solutions
skill lab - lab
Pg 216 # 1-9
Introduction to Acids
Chapter 5 review
and Bases
What is an acid?
Reactions that
What is a base?
generate acids and
- Pg 237 5,7
bases
- Pg 239 1-3
- Pg 242 4,5,7
- Pg 243 9 – 11
-Acids naming
worksheet
Calculation of PH
- Chapter 5
andPOH
skills practice
Auto-ionization of
worksheet
Week 13
Nov 24 - 28
Quantitative
Chemical Change
Week 14
Dec 1 – 5
Quantitative
Chemical Change
Week 15
Dec 8 - 12
Quantitative
chemical Change
Week 16
Dec 15 - 19
Quantitative
chemical change
Gravimetric
stoichiometry
water
Acid base reactions
Buffer solutions
Monoprotic and
polyprotic acids and
bases
-
Review / review
-Unit Test
solutions Acids
and bases
Review Science
10- bonding
Gases
Cumulative
booklet finish all
questions
Midterm test
Decomposition of
baking soda
Terms procedures
types of stoich and
examples
Nelson 7.2
Indicators lab
Pg 247 1-4
Pg 251 1,2
Pg 253 1-8
Pg 255 4,5
Pg 257 7-9
Pg 259 3-10
Pg 265 unit 3
review – 1-17,
Part 2-odds
Pg 276-283 Net
ionic equations
Questions 1-6
Pg 272 1-10
Pg 284 -290
#8-14
Pre-lab calculation
Labe write up
from 7.2
Percent yield and
Read pg 290 -293
difference
Lab exercises 7A,B
Gas stoichiometry
293 #8-10
SATP – STP
Quiz net ionic
Gas Stoich PV = nRT
equations
Solution Stoichiometry Worksheets Gas
Stoichiometry and
solution
stoichiometry
Chapter 7 review
Chapter 7 review
pg 309 1-14, 2031, 33
Chapter 7 test
This schedule unfortunately does not finish the course. The regular teacher will still need to cover
titrations and qualitative analysis – chapter 8 issues + review for course final.