SCH4U EXAM Advice
PART A
PART B
PART C
KNOWLDEGE AND UNDERSTANDING
Multiple Choice (5 or 6 from each unit)
APPLICATION / THINKING & INQUIRY
Short Answer
15 questions X 3 marks each (3 from each unit)
COMMUNICATION
Essay
2 questions X 6 marks each (1 mandatory, 1 choice)
29
45
12
86
WEIGHTING:
30 %
Provided: Periodic Table; Polyatomic ion chart; Solubility Chart; Redox table
Bring: Calculator, pen, pencil
All 5 units covered.
General Exam Advice:
1. To study: go over all notes, quizzes, labs and unit tests. Make all corrections.
2. The exam you care about the most, study the most for.
3. The credit that you need most, study for that course’s exam the most.
4. Make study notes for each unit. Start off by making several pages of notes for each unit. Condense to a one page study note for each unit.
5. Do not sit there and read over notes – be active!!! Write, summarize, cover up the answer and solve the problem without peeking until you’re
completely done the problem.
Course Specific Advice:
Unit
Atomic Theory
Areas of focus
- Explain the experimental observations and inferences made by Rutherford and Bohr in developing the planetary
model of the hydrogen atom. (3.1, 3.2, 3.3, 3.4)
- Describe the contributions of Planck, Bohr, Sommerfeld, de Broglie, Einstein, Schrödinger, and
Heisenberg to the development of the quantum mechanical model. (3.3, 3.4, 3.5, 3.6, 3.7)
- Describe the quantum mechanical model of the atom.(3.5, 3.6, 3.7)
- Write electron configurations for elements in the periodic table, using the Pauli exclusion principle and Hund’s
rule. (3.5, 3.6)
- Describe some applications of principles relating to atomic structure in analytical chemistry and medical diagnosis
(NMR/MRI)
- Explain how the Valence-Shell-Electron-Pair-Repulsion (VSEPR) model can be used to predict
molecular shape. (4.3)
- Predict molecular shape for simple molecules and ions, using the VSEPR model. (4.3)
- Predict the polarity of various substances, using molecular shape and electronegativity values of the elements of
the substances. (4.4)
- Explain how the properties of a solid and liquid depend on the nature of the particles present and the types of
forces between them. (4.4, 4.5, 4.6)
- Predict the type of solid (ionic, molecular, covalent network, or metallic) formed by a substance, and describe its
properties. (4.6)
Organic
- Distinguish among the different classes of organic compounds, including alcohols (1.5), aldehydes (1.6), ketones
(1.6), carboxylic acids (1.7), esters (1.7), ethers (1.5), amines (1.8), and amides (1.8), by name and by structural
formula.
- Describe some physical properties of the classes of organic compounds in terms of solubility in different solvents,
molecular polarity, odour, and melting and boiling points. (all sections)
- Describe different types of organic reactions, such as substitution (1.3, 1.4, 1.8), addition (1.4), elimination (1.4,
1.5), oxidation (1.5, 1.6, 1.7), esterification (1.7), and hydrolysis (1.7, 1.8), and predict and correctly name their
products. (1.3, 1.4, 1.5, 1.6, 1.7, 1.8)
- Describe some physical properties of the classes of organic compounds in terms of solubility in different solvents,
molecular polarity, odour, and melting and boiling points. (2.1, 2.5)
- Demonstrate an understanding of the processes of addition and condensation polymerization. (2.1, 2.2, 2.4, 2.5)
- Describe a variety of organic compounds present in living organisms, and explain their importance to those
organisms. (2.4, 2.5, 2.6, 2.7)
2
Energy and Rates
• Determine enthalpy of reaction using a calorimeter, and use the evidence obtained to calculate the
enthalpy change for a reaction. (5.1, 5.2, 5.3, 5.4)
• Compare the energy changes resulting from physical change, chemical reactions, and nuclear reactions (fission and
fusion). (5.1, 5.6)
• Describe examples of technologies that depend on exothermic or endothermic changes. (5.1, 5.6)
• Analyze simple potential energy diagrams of chemical reactions. (5.3)
• Write thermochemical equations, expressing the energy change as a delta H value or as a heat term in the equation.
(5.3)
• Explain Hess’s law, using examples. (5.4)
• Apply Hess’s law to solve problems, including problems that involve evidence obtained through
experimentation.(5.4, 5.5)
• Calculate enthalpies of reaction, using tabulated standard enthalpies of formation. (5.5)
• Describe, with the aid of a graph, the rate of reaction as a function of the change of concentration of a reactant or
product with respect to time. (6.1)
• Explain, using collision theory and potential energy diagrams, how factors such as temperature, surface area,
nature of reactants, catalysts, and concentration control the rate of chemical reactions. (6.2, 6.5)
• Express the rate of reaction as a rate law equation. (6.3)
• Determine a rate of reaction experimentally, and measure the effect on rate of temperature, initial concentration,
and catalysis. (6.3)
• Analyze simple potential energy diagrams of chemical reactions. (6.4)
Equilibrium
• Illustrate the concept of dynamic equilibrium with reference to systems such as liquid-vapour equilibrium, weak
electrolytes in solution, and chemical reactions. (7.1)
• Demonstrate an understanding of the law of chemical equilibrium as it applies to the concentrations of the
reactants and products at equilibrium. (7.2)
• Demonstrate an understanding of Le Châtelier’s Principle; and apply the principle to predict how
factors (such as changes in volume, pressure, concentration, or temperature) affect a chemical system at equilibrium,
and confirm your predictions through experimentation. (7.3)
• Explain how equilibrium principles may be applied to optimize the production of industrial chemicals. (7.4)
• Solve equilibrium problems involving concentrations of reactants and products, and K (7.2) and Ksp (7.5).
• Describe, using the concept of equilibrium, the behaviour of ionic solutes in solutions that are unsaturated,
saturated, and supersaturated. (7.6)
• Define constant expressions, such as Ksp. (7.6)
• Carry out experiments and calculations to determine equilibrium constants. (7.6)
• Predict the formation of precipitates by using the solubility product constant, Ksp. (7.6)
• Identify, in qualitative terms, entropy changes associated with chemical and physical processes. (7.7)
• Describe the tendency of reactions to achieve minimum energy and maximum entropy. (7.7)
• Define constant expressions, such as Kw (8.1), Ka (8.2), and Kb (8.2).
• Compare strong and weak acids and bases using the concept of equilibrium. (8.1, 8.2)
• Solve equilibrium problems involving concentrations of reactants and products and the following quantities: Ka,
Kb, pH, pOH. (8.2, 8.3, 8.4)
Electrochemistry
• Demonstrate an understanding of oxidation and reduction in terms of the transfer of electrons or
change in oxidation number. (9.1)
• Write balanced chemical equations for redox reactions using half-reaction equations. (9.1, 9.2, 9.3) – BALANCE
REDOX!
• Predict the spontaneity of redox reactions and cell potentials using a table of half-cell reduction
potentials. (9.3)
• Describe galvanic cells in terms of oxidation and reduction half-cells and electric potential differences.(9.5)
• Describe the function of the hydrogen reference half cell in assigning reduction potential values. (9.5)
• Determine oxidation and reduction half-cell reactions, current and ion flow, electrode polarity and cell potentials of
typical galvanic cells. (9.5)
• Predict the spontaneity of redox reactions and cell potentials using a table of half-cell reduction potentials. (10.1)
- distinguish primary cell from secondary cell
3
SCH4U EXAM REVIEW
SEC. 1
1. Propane can be used as starting material for either propanal or propanone. Explain how this is possible.
2. Make a summary table of organic functional groups using the heading: Name, Functional Group, IUPAC Suffix.
3. The space filling model for rayon is given below.
The large dark atoms are carbon, the medium-sized gray atoms are oxygen, and the small white atoms are hydrogen.
Identify the monomer needed to produce rayon, and explain why it is a tough fabric.
4. The structures of three amino acids are given below. List all the simple proteins which can be produced from these
amino acids.
5. Draw the electron configuration for oxygen and explain how you use Hund's Rule and the Pauli Exclusion principle to
do it.
6. Draw the electron configuration for the chloride, Cl1-, ion.
7. How are spectrophotometers used to identify gases in the sun?
8. Why was de Broglie's contribution to science so important to the development of Quantum Theory?
9. Explain why CH3F is a polar molecule while CF4 is not.
10. What would be the shape of NO3-1? Explain your reasoning.
11. Is ammonia a polar molecule? Explain your answer using a diagram.
12. In your own words, define what is meant by the rate of a reaction. Give an example of how a rate of reaction can be
determined.
13. Describe how the rate of a reaction such as,
, could be increased.
14. Differentiate between an effective collision and an ineffective collision.
15. Consider the equilibrium below:
If 1.6 mol of HI was placed in a 1.0 L container and allowed to reach equilibrium, what would the equilibrium
concentrations be for H2(g), I2(g) and HI(g) if the Ke = 36?
H2(g) + I2(g) <=====> 2HI(g)
16. Write an equilibrium expression showing how the fluoride ion acts like a base. Be sure to identify the conjugate acidbase partners in the equilibrium.
4
17. What is the oxidation number of iodine in
?
18. What is the algebraic sum of the oxidation numbers of all the atoms in aluminum sulfate,
?
19. Will a reaction take place if a strip of zinc metal is placed in a solution of lead (II) sulfate? Justify your answer.
SEC. 2
20. Calculate the enthalpy change,
for the vaporization of 200 g of methanol.
21. Given the following information, determine the molar heat of vapourization of water.
H2(g) + 1/2 O2(g)  H2O(g) + 242.8 kJ
H2(g) + 1/2 O2(g)  H2O(l) + 286.9 kJ
22. a. Calculate Ho for the following reaction.
b. State whether the reaction is exothermic or endothermic.
c. Rewrite the equation as a thermochemical equation to include the heat term.
d. Indicate whether the products have a greater or smaller enthalpy than the reactants.
23. The following data was obtained in a chemical kinetics experiment in which the concentration was measured every 10
s.
Time (s)
Concentration (mol/L)
0
0
10
0.13
20
0.25
30
0.34
40
0.40
50
0.43
60
0.45
70
0.46
80
0.46
90
0.47
a. Determine the average rate of formation of product during the period from 20 s to 70 s.
b. Plot a properly labelled graph of Concentration vs. Time.
c. Use the graph to estimate the average rate of formation during the same period as in a.
24. If the concentration of CO32- is found to be 1.3  10-4 mol/L in a saturated solution of Ag2CO3, what is the Ksp of
Ag2CO3?
5
25. If 45 mL of a 0.45 mol/L solution of AgNO3 was mixed with 85 mL of a 1.35  10-2 mol/L solution of NaCl, would a
precipitate form? Calculate the ion product for the potential precipitate. The Ksp of AgCl(s) is 1.8  10-10.
26. What is the concentration of a weak base if its Kb = 1.4  10-11 and its pH = 8.75?
27. What is the percent ionization of a 0.48 mol/L weak acid if its Ka = 1.4  10-9?
28. What voltage is necessary to force the following electrolysis reaction to occur?
SEC. 3
29. Balance the following equation by the oxidation number method and identify the oxidizing and reducing agents:
30. Balance the following redox reaction by the oxidation number method:
6
7
SCH4U EXAM REVIEW
Answer Section
SEC. 1
1.
2.
Name
Functional Group
alcohol
IUPAC Suffix
-ol
ether
-ether
aldehyde
-al
ketone
-one
carboxylic acid
-oic acid
ester
-oate
amine
-amine
amide
-amide
3. glucose is the monomer:
hydrogen bonding between adjacent polymer molecules increases strength
4.
- gyl-ala
- gly-ser
- ala-ser
5.
1s22s22p4
- gly-ala-ser
- gyl-ser-ala
- ala-gly-ser
8
Pauli exclusion principle states that no two electrons can have the same four quantum numbers, therefore
each orbital can hold only two electrons with opposite spins. Hund's rule says that the electrons in orbitals
with the same energy are half filled first before more are added. Also, the electrons in those half filled
orbitals must have the same spin.
6.
7.
8.
9.
10.
11.

3p

3s

2p

2s

1s
By noting the line spectra produced by sun light and explaining its colours by comparing it to the line spectra
of different gases.
- formalized the duality of matter
- allowed scientists to envision electrons as wave/particles
Both species have a tetrahedral structure. This shape is symmetrical in all directions. Therefore, even though
the CF4 molecule contains four polar bonds, there is no positive end and negative end; therefore the molecule
itself is not polar. However, the CH3F molecule contains only one C-F bond. The F draws electrons from the
rest of the molecule as well, forming a partial negative charge at its end, and leaving the hydrogens with
partial positive bonds.
It would be a trigonal planar molecule, because nitrogen would have three oxygen atoms and no unbonded
pairs of electrons around it. This configuration would get the bonded electrons as far apart as possible.
Ammonia would be polar because it has a lone pair that pushes the hydrogens down, forming an asymmetric
molecule. Also, nitrogen is highly electronegative, so each N-H bond is polar.
12. 13. 14. -
answer should indicate a rate of change in the reactants or products over a given period of time
rate can be determined by measuring the change in mass, colour, conductivity, volume, pressure, etc.
since concentration is a factor that affects the rate of reaction, by increasing the concentration of the acid,
the reaction rate would be increased
temperature is another factor that affects the rate of reaction, by increasing the temperature of the
reaction, the reaction should proceed at a faster rate
since surface area of a solid would affect the rate of a reaction, by powdering the magnesium, the rate of
the reaction would be increased
perhaps a catalyst could be added that could speed up the reaction (unknown)
perhaps using a different acid and/or metal (although this would change the reaction though)
in an effective collision, the particles collide with sufficient energy and correct orientation for a reaction
to occur, the formation of products occurs from this collision
in an ineffective collision, the particles collide without sufficient energy and/or correct orientation for the
reaction to occur, the reactants are unchanged by this collision
15.
mol/L
H2(g) + I2(g) <=====> 2HI(g)
initial
1.6
shift
–2


@E
1.6  2


Ke = (1.6  2)2 / (2) = 36 (take the square root of both sides)
1.6  2 /  = 6
6 = 1.6 –2
8 = 1.6
9
 = 0.20 mol/L = [H2] = [I2]; [HI] = 1.6  2(0.20) = 1.2 mol/L
16.
F1-(aq) + H2O(l)
base-1 acid-2
17. +5
18. zero
<=====>
HF(aq) + OH1-(aq)
acid-1
base-2
19. Oxidation half-reaction:
Reduction half-reaction:
Overall reaction:
for the overall reaction is positive. Thus a zinc strip will react spontaneously with lead (II) sulfate
solution.
SEC. 2
20.
Find the number of moles of methanol
Then find the enthalpy change
Since the methanol vaporizes by absorbing heat, the enthalpy change is +244.8 kJ.
21. Show the addition of equations to get and thus appropriate
values
H2O(l)  H2O(g)
H2(g) + 1/2 O2(g)  H2O(g)
H2O(l)  H2(g) + 1/2 O2(g)
H2O(l)  H2O(g)
The molar heat of vaporization of water is 44.1 kJ/mol based on the given information.
22. a.
The heat of reaction is –907 kJ/mol of Benzene
b. Since the sign of the heat of reaction is negative, the reaction is exothermic.
10
c.
d. Since the reaction is exothermic, the products would have a smaller enthalpy than the reactants, since
energy is given off to the surroundings.
23. a.
The rate of production during the 50s time period is 0.0042 mol/(L·s)
b. The graph should look somewhat like the following graph.
c. These answers will vary depending on the students graph. A secant should be drawn on the graph from
20s to 70 s and the slope of the line calculated using rise over run.
11
24.
Ag2CO3(s) <====>
@E
2Ag1+(aq)
2 * 1.3  10-4
+
CO32-(aq)
1.3  10-4
Ksp = ( 2 * 1.3  10-4 )2  1.3  10-4
Ksp = 8.8  10-12
25. AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
[Ag1+] = (0.45 mol/L)(0.045 L) / (0.045 L + 0.085) L
[Cl1-] = (1.35  10-2 mol/L)(0.085 L) / (0.045 L + 0.085) L
ion product AgCl is [Ag1+][Cl1-] = 1.4  10-3 > Ksp, yes a precipitate forms
26.
XOH(aq) <=====> X1+(aq) + OH1-(aq)
x
10-pOH
10-pOH
pOH = 14 – pH = 14 – 8.75 = 5.25
[OH1-] = 10-pOH = 10-5.25
[OH1-] = 5.6  10-6 mol/L
(5.6  10-6)2 /  = 1.4  10-11 mol/L
x = (5.6  10-6)2 / 1.4  10-11 = 2.3 mol/L
27.
HY(aq) <====> H1+(aq) + Y1-(aq)
0.48
-x
x
x
x
x
0.48  x
initial
shift
@E
x2 / 0.48 = 1.4  10-9
x = (1.4  10-9  0.48)0.5
pH = – log x
pH = 4.59,
% Ionization = 100(x / 0.48) = 5.4  10-3 %
28.
The anode reaction:
The cathode reaction:
The cell voltage:
More than 1.60 V must be externally applied to make this reaction proceed.
SEC. 3
29.
-1
+6 2
+1
0
+1 2
+1 2
The change in the oxidation number of the oxidized element I = (0) – (–1) = +1
The change in the oxidation number of the reduced element S = (–2) – (+6) = –8
The lowest common multiple of 1 and 8 is 8. Thus in the balanced reaction eight iodine atoms are needed for
every sulfur atom. The increase and decrease in oxidation numbers will then be eight for both. The partially
balanced equation is:
The remaining coefficients can be balanced by inspection and the balanced equation is:
12
The oxidizing agent is:
. The reducing agent is
.
30.
+2
2
3 +1
0
+1 2
0
The change in the oxidation number of the oxidized element N = (0) (–3) = +3
The change in the oxidation number of the reduced element
= (0) (+2) = –2
The lowest common multiple of 3 and 2 is 6. Thus in the balanced reaction two nitrogen atoms are needed
for every three copper atoms. The increase and decrease in oxidation numbers will then be six for both. The
partially balanced equation is:
The remaining coefficients can be balanced by inspection and the balanced equation is: