SCH 4U – GRADE 12 CHEMISTRY FINAL EXAM
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Part 1: Multiple Choice
[40 MARKS]
Fill in the correct answer in the blank beside each question.
_____ 1. An amine is characterized by what functional group?
A. -CO2CH3
B. -NH2
C. -CO2H.
D. –CHO
E. –CONH2
_____ 2. Why does the boiling point of an alkane increase as its chain length increases?
A. There is more hydrogen bonding.
B. There are more hydrogen atoms available for hydrogen bonds to form.
C. The number of dipole-dipole interactions increases.
D. The strength of the dispersion forces increases with increased molecular size.
E. Heavier molecules cannot float on the surface of water that well.
_____ 3. Which substance is not a structural isomer of hexyne?
A. 2-hexyne
B. 3-hexyne
C. 2,2-dimethylpentyne
D. 4-methyl-1-pentyne
E. 2,3-dimethylbutadiene
_____ 4. Which of the following is a secondary alcohol?
A.
B.
D. CH3OH
E. CH3CH2OH
_____ 5. What is the correct name for the compound below?
A. Chloropropylbenzene
B. 3-phenyl-1-chloropropyne
C. 1-chloro-3-phenyl-3-propene
D. 3-chloro-1-phenyl-1-propene
E. 1-chloro-3-phenyl-2-propene
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C.
_____ 6. Markonikov’s Rule states that when an alkene or alkyne reacts with either a
hydrohalide or with water that:
A. the carbon that already had the most H atoms receives the H atom
B. the H atoms are lost as H2 gas
C. Markonikov’s Rule does not apply to this situation
D. the carbon that has the fewest H atoms receives the H atom
E. the H atoms combine with O to form water
_____ 7. Which type of reaction will an alkene NOT undergo?
A. Addition
B. Polymerization
C. Oxidation
D. Dehydration
E. Hydration
_____ 8. The synthesis sequence shown here is best described as which of the following?
1.
2.
3.
A. (1) Dehydration; (2) halogenation; (3) hydrogenation
B. (1) Hydrogenation; (2) dehydration; (3) halogenation
C. (1) Hydrogenation; (2) halogenation; (3) dehydration
D. (1) Halogenation; (2) hydrogenation; (3) dehydration
E. (1) Dehydration; (2) hydrogenation; (3) halogenation
_____ 9. The 3p atomic orbital has the shape of:
A. a sphere
B. a donut
C. a dumb-bell
D. two perpendicular dumb-bells
E. an egg
_____ 10. Why do energy levels exist in atoms?
A. electrons are negatively charged
B. electrons are attracted to certain numbers of neutrons
C. electrons are able to possess any range of energy
D. electrons will only display certain colours
E. electrons are only able to possess quanta of energy
_____ 11. "A region of space in which there is a high probability of finding an electron"
is the definition of
A. orbital
B. absorption spectrum
C. quantum
D. photon
E. dipole
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_____ 12. Which of the following is the electron configuration for the valence shell of
oxygen?
A.
B.
C.
D.
E.
_____ 13. Which of the following is the electron configuration for magnesium?
A. 1s22s22p8
B. 1s32s32p43s2
C. 1s22s22p73s1
D. 1s32s32p33s2
E. 1s22s22p63s2
_____ 14. Which of the following are properties of alkali metals?
I.
II.
III.
IV.
They have one valence electron.
They have high first ionization energies.
They are very reactive.
Their most easily removed electron is in an s orbital.
A. I and II only
B. I and IV only
C. I, III and IV only
D. III only
E. I only
_____ 15. Four pairs of electrons surrounding a central atom will be arranged?
A. Pyramidally
B. Spherically
C. Tetrahedrally
D. Linearly
E. Octahedrally
_____ 16. Which of the molecules, CO2, H2O, NH3, and BF3, will be polar?
A. CO2, NH3 and BF3
B. H2O and NH3
C. H2O and BF3
D. CO2, H2O and NH3
E. CO2 and BF3
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_____ 17. What would be the shape of a molecule containing a central atom attached to
two other atoms with one lone pair of electrons?
A. Linear.
B. Bent.
C. Trigonal Planar.
D. Tetrahedral.
E. Trigonal bipyramidal
_____ 18. Metals can be rolled into sheets and stamped into various forms. In contrast,
diamond is very hard and brittle. Which explanation for these different properties is
correct?
A. Metals have semi-ionic bonds, whereas diamond has covalent bonds.
B. The electrons that surround a metal atom are free to move through the metal. The
bonding electron pairs in a diamond are held tightly between two carbon atoms in an
overall tetrahedral pattern.
C. The electrons of a metal are held more tightly to the parent atom than the electrons
of carbon. Hence, the bonds in a metal are stronger than the bonds in diamond.
D. Diamond has strong double bonds between carbon atoms. Metal bonds are
normally single covalent bonds, which bend easily.
E. Metals are made of metal atoms, whereas diamond is made of non-metal carbon
atoms.
_____ 19. An exothermic reaction is one where:
A. heat is transferred from the surroundings into a system
B. heat is transferred from a system into the surroundings
C. kinetic energy is transformed into potential energy
D. there is no transfer of heat
E. None of the above.
_____ 20. What is the definition of temperature of a substance?
A. the total heat content of a substance.
B. the speed of the fastest particles in the substance.
C. the speed of the slowest particles in the substance, subtracted from the speed of the
fastest particles in the substance.
D. the heat capacity of the substance times its mass.
E. the average kinetic energy of a system.
_____ 21. In a calorimeter, a 1.0 g sample of magnesium is burned to form MgO. In
doing so, 25.5 kJ of energy are released. What is the Heat of Combustion in kJ/mol of
magnesium?
A. 306.2.
B. 1.54 x 1025.
C. 0.0392.
D. 25.5.
E. 620
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_____ 22. Which statement concerning the accompanying diagram is true?
A. ΔH is positive
B. the system is endothermic
C. the system releases heat to the surroundings
D. the heat content of the reactants is less than the heat content of the products
E. the enthalpy of the products is greater than the enthalpy of reactants
_____ 23. Given the following thermochemical data:
1. C2H2(g) + 5/2 O2(g)  2CO2(g) + H2O(l)
2. C2H6(g) + 7/2 O2(g)  2CO2(g) + 3H2O(l)
3. H2(g) + 1/2 O2(g)  H2O(l)
ΔH = –1.30 x 103 kJ
ΔH = –1.56 x 103 kJ
ΔH = –2.86 x 102 kJ
What is ΔH for the following reaction?
C2H2(g) + 2H2(g)  C2H6(g)
A. –2.60 x 101 kJ
B. –3.12 x 102 kJ
C. –5.72 x 102 kJ
D. –5.46 x 103 kJ
E. 2.60 x 101 kJ
_____ 24. Consider the following equation for the combustion of hydrogen:
H2(g) + 1/2 O2(g)  H2O(g) + 243 kJ
In order to produce 1215 kJ of heat, how many grams of H2 must burn?
A. 12.0 g
B. 0.100 g
C. 10.0 g
D. 0.250 g
E. 8.00 g
_____ 25. The presence of a catalyst is thought to increase the rate of a reaction by:
A. changing the products that are formed in the reaction.
B. decreasing the enthalpy change of the reaction.
C. increasing the enthalpy change of the reaction.
D. decreasing the activation energy of the reaction.
E. increasing the activation energy of the reaction.
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_____ 26. If for the reaction ,
be
the rate law is determined to
, then
A. the order of the reaction is 0.
B. increasing the concentration of Y will have no effect on the rate.
C. increasing the concentration of X will have no effect on the rate.
D. increasing the concentration of Y will increase the rate of the reaction.
E. there is no way to determine the value of k.
_____ 27. In the following diagram, the letter which represents the position of the
activated complex is:
A. A
B. B
C. C
D. D
E. E
_____ 28. If a reaction can be broken down into a reaction mechanism, then the steps of
the reaction mechanism are known as:
A. stages of reaction
B. activated complexes
C. reaction progress
D. elementary steps
E. primary equations
_____29. Consider the above reaction mechanism. The rate-determining step of this
reaction is
A. elementary step 1
B. elementary step 2
C. elementary step 3
D. elementary steps 2 and 3
E. impossible to tell from this information
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_____ 30. Consider the above reaction mechanism. The rate-law equation from this
reaction would be
A.
B.
C.
D.
E. impossible to tell from this information
_____ 31. For the equilibrium system below, which of the following would result in a
decrease in the quantity of PCl5(g)?
PCl3(g) + Cl2(g) <=====> PCl5(g) + 45 kJ
A. increasing temperature
B. adding some Cl2(g)
C. decreasing temperature
D. decreasing the size of the container
E. injecting some He gas
_____ 32. A concentrated weak acid is best described as which of the following?
A. a solution with a low pH
B. a solution where the concentration of undissociated acid particles is low compared
to the concentration of hydronium ions
C. a solution where the concentration of hydronium ions is large compared to the
concentration of undissociated acid particles
D. a solution with a high pH
E. a solution where the concentration of undissociated acid particles is high and the
relative quantity of hydronium ions is small
_____ 33. The Ksp for a substance A2B is 1.2 x 10-12. The [B] must be which of the
following?
A. 4.0 x 10-11
B. 3.6 x 10-12
C. 6.7 x 10-5
D. 1.3 x 10-4
E. 2.6 x 10-4
_____ 34. Which condition represents a spontaneous reaction?
A. Negative change in enthalpy and negative change in entropy.
B. Negative change in enthalpy and positive change in entropy.
C. Positive change in enthalpy and negative change in entropy.
D. Positive change in enthalpy and positive change in entropy.
E. All of the above
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_____ 35. In the reaction
+
+
Which of the following statements is correct?
A.
is the reducing agent, and
is the oxidizing agent
B.
is the reducing agent, and
is the oxidizing agent
C.
is the reducing agent, and
D.
is the reducing agent, and
E.
is the reducing agent, and
+
+
is the oxidizing agent
is the oxidizing agent
is the oxidizing agent
_____ 36. Which of the following statements are true for the redox reaction?
Cr2O72-(aq) + 14H+(aq) + 6Br -(aq)  2Cr3+(aq) + 3Br2(aq) + 7H2O
I.
II.
III.
IV.
V.
the oxidation number of Cr changes from +7 to +3
Cr2O72- is the oxidizing agent
Br - is oxidized to Br2
H+ is the reducing agent
the oxidation number of H+ changes from +1 to 0
A. II and IV
B. I and III
C. II and III
D. IV and V
E. III and V
_____ 37. The following two half-reactions are involved in a galvanic cell:
At standard conditions, what species are produced at each electrode?
A.
is produced at the cathode and
at the anode
B.
is produced at the cathode and
at the anode
C.
is produced at the anode and
at the cathode
D.
is produced at the cathode and
E.
is produced at the cathode and
at the anode
at the anode
_____ 38. Which of the following equations does NOT represent an oxidation-reduction
reaction?
A.
+
B.
+
C.
+
D.
+
E.
+
+
+
+
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_____ 39. The function of the salt bridge in a galvanic cell is:
A. to act as a nonelectrolyte
B. to provide an external circuit for the flow of electrons from one half-cell to the
other
C. to provide a path for the migration of ions from one half-cell to the other
D. to allow the solution in each half-cell to become electrically charged
E. to provide a path for the flow of electrons internally from one half-cell to the
other
_____ 40. The oxidation number of chromium (
A. + 2
B. + 4
C. + 6
D. + 7
E. - 2
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) in the dichromate ion (
) is
Part II: Short Answers / Calculations
[60 MARKS]
1. Prepare Ethyl Benzoate from Benzyl Alcohol and Ethene as your only sources of
carbon:
[5 MARKS]
2. Define the term "orbital." Describe the shape of each type of orbital. State how many
of each type are present, and in which energy levels they are located.
[3 MARKS]
3. Use your knowledge of equilibrium to explain the difference between a strong base
and a weak base.
[2 MARKS]
4. What is the oxidation number of nitrogen in each compound?
a) Li3N
[2 MARKS]
b) N2
c) N2H4
d) NH4NO3
5. What would be the predicted shape of NO3-1 (Include the Lewis structure and VSEPR
diagram). Is this molecule polar or non-polar. Explain.
[3.5 MARKS]
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6. Complete the following table.
[4 MARKS]
7. Fill in the following outline of a periodic table to show the four energy sublevels
(s, p, d, and f) blocks.
[2 MARKS]
8. A hypothetical molecule has the formula XY4. The X-Y bond is polar covalent, but the
molecule is not polar. Which molecular shapes are possible and impossible for the
molecule? Explain why, for each shape.
[4 MARKS]
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9. Balance the following redox reaction that takes place in a basic solution and determine
which species is oxidized and which is reduced:
[3 MARKS]
NO3-1 (aq) + Al (s)  NH3(aq) + Al(OH)4-1(aq)
10. Solve the ∆Hr for the following balanced equation using Hess’s Law and the given data:
C3H8 (g) + 5O2 (g)  3CO2 (g) + 4H2O (g)
[2 MARKS]
C3H8(g)  3C(s) + 4H2(g)
C(s) + O2(g)  CO2(g)
H2O(l)  H2(g) + 1/2 O2(g)
∆H1 = +104 kJ
∆H2 = -393.5 kJ
∆H3 = +285.5 kJ
11. What is the pH of a 1.24 mol/L solution of HCN(aq) if its Ka = 6.2 x 10-10?
[1 MARK]
12. Write the electron configuration and energy level diagram for vanadium 2+ ion.
Label completely.
[3 MARKS]
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13. Use the following data to calculate the reaction rate law and the rate law constant for
the system. (Include units!)
[4 MARKS]
Initial Rate of Reaction (mol/(L·s))
H2 (mol/L)
NO (mol/L)
Experiment
0.002
0.004
0.001
1
0.008
0.004
0.002
2
0.018
0.004
0.003
3
0.008
0.001
0.004
4
0.016
0.002
0.004
5
0.024
0.003
0.004
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14. In a calorimeter, a 1.0g sample of magnesium is burned to form MgO. In doing so,
100g of water in the calorimeter had a temperature increase of 60oC. What is the molar
heat of combustion for magnesium in kJ/mol? (cwater = 4.18J/goC)
[1.5 MARKS]
15. Use half-cell potentials to predict whether the following reactions are spontaneous or
not.
[2 MARKS]
2+
2+
3+
a) Cu (aq) + Fe (aq) → Fe (aq) + Cu(s)
b) Ni(s) + Ag+ (aq) → Ni2+(aq) + Ag (s)
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16. Draw the products and name them for the following reactions.
[1 / 2 / 1 MARKS]
a)
b)
c)
17. An aqueous 1M solution of NiSO4 was electrolyzed using inert electrodes. What
substance was produced at each electrode? What is the minimum voltage that must be
applied to the cell?
[2.5 MARKS]
18. In a titration of 20.00 mL of 0.300 mol/L HC2H3O2(aq) with standardized 0.300 mol/L
NaOH(aq), what is the pH of the solution:
(a) before titration begins; Ka of HC2H3O2 = 1.8 x 10-5
[2 MARKS]
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(b) during titration but before the equivalence point (10.00 mL of 0.300 mol/L NaOH(aq)
added);
[3 MARKS]
(c) at the equivalence point (20.00 mL of 0.300 mol/L NaOH(aq) added);
[3 MARKS]
(d) beyond the equivalence point (40.00 mL of 0.300 mol/L NaOH(aq) added) [1.5 MARKS]
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e) Draw a sketch of the titration curve
[2 MARKS]
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