Gr.12 Chemistry Exam Review Guide June 2015
Use old tests : these are the best and most efficient way to study. Do the corrections, and then
go to your notes to fill in your gaps in knowledge. You may also wish to make yourself study
notes from the guide below. I have included the learning outcomes that may be examined for
each unit, along with question types for the first couple of units.
Website: Worksheets and detailed notes for most topics are on the website. Also, don’t forget
to check out the videos when you’re getting a hand cramp writing down total ionic equations
and the like.
A Periodic Table, relative acid/base strength chart, oxidation number rules, and solubility rules
will be provided. ****BRING A CALCULATOR*****
Unit 1 – Aqueous Solutions
1) Explain observed examples of solubility and precipitation at the molecular and symbolic levels.
2) Perform a lab to develop a set of solubility rules. Use a table of solubility rules to predict the
formation of a precipitate.
3) Write balanced neutralization reactions involving strong acids and bases.
4) Calculate the concentration or volume of an acid or base from the concentration and volume of
an acid or a base required for neutralization.
5) Define oxidation and reduction. Include: gain and loss of electrons, oxidizing agent, reducing
agent.
6) Determine the oxidation numbers for atoms in compounds and ions.
7) Identify and describe reactions as redox or non-redox. Include: oxidizing agent, reducing agent,
oxidized substance, and reduced substance.
8) Balance oxidation-reduction reactions using redox methods.
Question types:
1.
2.
3.
4.
5.
6.
7.
Draw pictures to represent what happens at the molecular level of the following reaction:
Define precipitate. Which compound(s) in question #1 is/are precipitate(s)?
For the following questions, complete the equations, adding in (s), (l), (aq), or (g) where
appropriate. Use the solubility table provided.
For the reaction _________________ write the:
a) Molecular equation – Make sure to include (aq) and (s). Use the solubility table provided.
b) Total Ionic Equation
c) Net Ionic Equation
d) What’s a spectator ion?
e) List the spectator ions in this reaction.
Compare the traditional definitions of oxidation and reduction to the modern definition.
In the following reactions, label the oxidizing agent and the reducing agent.
For these example reactions, determine: the substance being oxidized, the substance being
reduced, the oxidizing agent, and the reducing agent.
Oxidation Number Rules:
You will be given these exact rules on your exam, and expected to determine the oxidation
numbers for a specific element in compounds:
a. The oxidation number of any free atom (or multiple of itself) is 0.
b. An ion’s (or complex ion’s) oxidation number is its charge.
c. In a compound the sum of all of the oxidation numbers of each part must equal the total
charge of that compound or complex ion.
d. Hydrogen’s oxidation number is +1, except in metal hydrides where H is the where it is 1.
e. The oxidation number of oxygen is -2, except in peroxides (ie. H2O2, Na2O2) where it is 1, and when in combination with fluorine (then O = +2)
8. Balance oxidation-reduction reactions using redox methods
Unit 2 – Kinetics
1)
2)
3)
4)
5)
6)
7)
8)
Formulate an operational definition of reaction rate. Include: examples of chemical reactions that occur
at different rates.
Identify variables used to monitor reaction rates (i.e. change per unit time, Δx/Δt). Examples: pressure,
temperature, pH, conductivity, colour
Calculate the average and instantaneous rate of a chemical reaction from a data table or a graph.
Relate the rate of formation of a product to the rate of disappearance of a reactant given experimental
rate data and reaction stoichiometry.
Use the Collision Theory to explain the factors which affect the rate of chemical reactions. Include:
Activation energy, orientation of molecules
Draw potential energy diagrams for endothermic and exothermic reactions. Include: relative rates, effects
of catalyst, heat of reaction (enthalpy change)
Explain the concept of a reaction mechanism. Include: rate determining step
Determine the rate law and order of a chemical reaction from experimental data. Include: reactions that
are zero, first or second order, rate versus concentration graphs.
For question types, you have booklets that had lots of these types of questions. You may also with to go
to check out: http://colgurchemistry.com/Chem12/chem12.htm for extra practice. Some of the
questions on these worksheets are a little different than what we did, but most are pretty much the
same.
Go to Unit 1: Reaction Kinetics
Worksheet 1-1 - Measuring Reaction Rates
Worksheet 1-2 - Potential Energy Diagrams
Worksheet 1-3 - Reaction Mechanisms (we did this one in class)
Unit 3 – Atomic Structure
1) Wavelength and Energy
a. Qualitatively describe the electromagnetic spectrum in terms of frequency, wavelength and
energy.
b. Explain why emission spectra are unique to each element
2) Historical Timeline
a. Outline the Development of the Quantum Mechanical Model of the atom
i. Dalton, Thompson, Rutherford, Bohr and Schrodinger
3) Electron Configurations
a. Write electron configurations for elements of the periodic table.
i. Full electron configurations atoms and ions
ii. Orbital configuration (spins)
iii. Noble gas simplification
b. 3 principles of electron fill
i. The Aufbau principle
ii. The Pauli Exclusion principle
iii. Hund’s rule
c. Identify
i. Valence electrons
ii. What element (based on configuration)
iii. Invalid configurations
4) Periodic Trends
a. Identify periodic trends among the properties of elements and relate to electron configuration.
Include: atomic radii, ionic radii, ionization energy, electronegativity
Question Types:
1. Explain that frequency and wavelength have an inverse relationship, while frequency and energy
have a direct relationship. Apply this to simple situations. (e.g. If wavelength is doubled, what
happens to frequency? If wavelength is halved, what happens to energy?)
2. Explain why each element has its own unique line spectra, in terms of energy levels.
3. Historical development of the Quantum Mechanical Model of the atom
a. Outline the major advances and flaws in the following models: in Dalton's, Thomson's,
Rutherford's and Bohr's How were the atomic models developed when no one had seen
the atom?
b. What flaws exist in Dalton's, Thomson's, Rutherford's and Bohr's model of the atom? Do
any flaws exist in the modern quantum model of the atom? Explain
c. Which of Dalton's principles were contradicted by the work of J.J.Thomson? Do any of
Dalton's principles still hold true completely today? If so, which ones hold true?
d. Describe Rutherford's model of the atom. Why was he so amazed by the results of his
experiment? What flaws exist in his model?
4. Write electron configurations with regular, spin and noble gas simplification notation– use
worksheets for practice
5. Pauli’s Exclusion Principle, Hunds’s Rule, Aufbau Principle
6. Identify periodic trends among the properties of elements and relate to electron configuration
(include atomic radii, ionic radii, ionization energy, electronegativity)
Unit 4 – Chemical Equilibrium Review
1) Equilibrium
a. Conditions necessary
b. Dynamic equilibrium
2) KEQ
a. write expressions
b. Ratio products: reactants
c. Write expressions
d. What values mean (>, < or = 1)
e. How temperature, pressure, catalyst, or concentration effect
3) Solve problems involving equilibrium constants.
a. Find Keq given concentrations at equilibrium
b. Given [I] and 1 [E] (ICE)
c. Given
4) Le Chatelier’s Principle
a. Use Le Chatelier’s Principle to predict shifts in equilibrium
i. Include: temperature changes, pressure / volume changes, changes in
reactant/product concentration, the addition of a catalyst,
b. Interpret concentration versus time graphs. Include: temperature changes,
concentration changes, addition of a catalyst
5) Ksp
a. Write solubility product (Ksp) expressions from balanced chemical equations
b. Solve problems involving Ksp.
i. Write dissociation equations
ii. Calculate Ksp given molar solubility
iii. Calculate molar solubility given Ksp
iv. Common ion problems
6) Enthalpy and entropy
a. Define
b. Identify which side of a reaction will be favors
c. Identify what occurs in a reaction based on its enthalpy/entropy
Review Questions:
Equilibrium (Keq) questions: (any boxes should be arrows – formatting problem, sorry)
1. What things are constant at equilibrium?
2. What can the equilibrium constant tell you about whether or not a reaction favours the products or
reactants?
3. Write the equilibrium constant expression for the following reactions: (Squares should be double
arrows – I can’t seem to fix the formatting issue.)
a) PCl5( g )
PCl3( g ) +Cl2( g )
b) 4NH3( g ) + 5O2( g )
4NO( g ) + 6H 2O( g )
5. Write the chemical reaction for the following equilibrium constant expression:
CH 4  H 2 S 
K eq 
4
 H 2  CS2 
2
6. For the chemical reaction H 2 + I 2
2HI , I start with 1 mol each of H2 and I2 in a 1 liter container.
Once equilibrium has been achieved, I have 1.56 mol HI. What is the value of the equilibrium constant
for this reaction?
7.for the chemical Reaction At a certain temperature the reaction:
CO(g) + H2O(g)
CO2(g) + H2(g)
If has a Keq = 0.400. Exactly 1.00 mol of each gas (not just the reactants!) was placed in a 100.0 L vessel
and the mixture was allowed to react. Find the concentration of each gas at equilibrium.
8. How does a catalyst change equilibrium situations?
9. What is Le Chatelier’s Principle?
11. In a closed container, the following reaction is allowed to come to equilibrium:
heat + NH 4 HS( g )
NH3( g ) + H 2 S( g )
How would the following stresses shift the equilibrium?
a) H2S added
b) temperature is increased
c) increase the size of the container
d) NH3 removed
e) NH4HS is added
f) write the Kc expression
12. Given the reaction: Cl2(g) + PCl3(g)
PCl5(g)
ΔH = -92.5 kJ
What part of the reaction does enthalpy favour? What part of the reaction does entropy favour? What
occurs during this reaction (ie. Equilibrium, completes, does not occur)
Solubility Product (Ksp) Questions:
1. Write dissociation reactions and the Ksp equations for when the following ionic compounds are
dissolved in water:
a) SrSO4 
b) Al2(CO3)3 
2. In water, the solubility of Al(OH)3 is 2.63x10-9 M. Find the Ksp value for Al(OH)3.
4. a) For MgF2, Ksp = 3.7x10-8. Find the solubility of MgF2 in water.
b) For Mg3(PO4)2, Ksp = 1x10-25. Find the solubility of Mg3(PO4)2 in water.
5. Calculate the solubility of MgF2 in a solution that contains 0.1 M MgCl2.
Unit 5 – Acids and Bases
1) Arrhenius definition & Bronsted Lowry definition of acids and bases
2) Identify what is a base and acid during acid/base equations
a. Identify conjugate acid/base pairs
3) Write balanced acid/base chemical equations for monoprotic and polyprotic acids
4) Solve problems involving pH, pOH, [H3O+] and [OH-]
5) Strong and weak acids and bases
a. Definition of both
b. Write the equilibrium expression (Ka or Kb) from a balanced chemical equation.
c. Use Ka or Kb to solve problems for pH, pOH, [H3O+] and [OH-].
d. Percentage dissociation
i. Be able to calculate for weak acid problem
Review Questions:
Use worksheets to help study for this section.