Gr ade 12 Chemistry

Final Practise Examination

Answer Key

e y
rK
For Marker’s Use Only
Name: ______________________________________

e
Student Number: _____________________________ Date: ___________________________

sw
Attending q Non-Attending q Final Mark _______ /100 = ________%

A n
Phone Number: ______________________________ Comments:

Address: ____________________________________

____________________________________________

____________________________________________

Instructions
You have a maximum of 3 hours to complete your final examination.
The final examination will be weighted as follows:
n Modules 4 to 6: 100%

The format of the examination will be as follows:

n Part A: Multiple Choice (23 x 1 = 23 marks)
n Part B: Fill-in-the-Blanks (19 x 1 = 19 marks)
n Part C: Short Answer (58 marks)

Include units with all answers as required.

You will need the following in order to complete this examination:

n writing utensils and eraser or correction fluid
n scrap paper
n a ruler
n a graphing or scientific calculator

The following resources are provided at the end of this examination:

n Periodic Table of Elements
n Alphabetical Listing of the Elements and Their Atomic Masses
n Table of Electronegativities
n Relative Strengths of Acids Table
n Solubility Chart
n Table of Standard Reduction Potentials with Values
n Names, Formulas, and Charges of Common Ions
n Common Ions

Final Practise Examination Answer Key 1 of 32

 Part A: Multiple Choice (23 marks total)

Use the answer sheet found at the end of this examination to answer the
multiple‑choice questions in this section. Shade in the circle that corresponds to
your answer. DO NOT circle your answers directly on the examination.
Equilibrium (7 marks)

1. Based on the equilibrium constants given, in which of these reactions are the products
most favoured over the reactants?
a. Keq = 0.002
b. Keq = 0.0
c. Keq = 3.5
d. Keq = 6.0 ×10-4
2. Which of the statements regarding the effect of reactant or product concentration
changes on equilibrium is false?
a. Adding a reactant always shifts a reversible reaction in the direction of reactants.
b. Removing a reactant always shifts a reversible reaction in the direction of
reactants.
c. Adding a product causes a reversible reaction to shift in the direction of the
formation of the reactants.
d. Removing a product shifts a reversible reaction in the direction of formation of
products.
3. Which of the following statements correctly describes the effect of an increase in
temperature on the position of equilibrium for an exothermic reaction?
a. The position of equilibrium does not change if you change the temperature.
b. According to Le Châtelier’s Principle, the position of equilibrium will move to the
left if the temperature is increased.
c. If you increase the temperature, the position of equilibrium will move in such a
way as to increase the temperature again.
d. Equilibrium will shift to favour the reaction which releases heat in this context.
4. Which explanation best describes why a catalyst does not affect the position of
equilibrium?
a. A catalyst will increase the activation energy in both directions; therefore, no shift
in the position of equilibrium will result.
b. The position of equilibrium is not affected by adding a catalyst since a catalyst
speeds up both the forward and back reactions by exactly the same amount.
c. Equilibrium constants are not affected by adding a catalyst since no shift in the
position of equilibrium occurs.
d. Not all chemical reactions respond to catalysts.

2 of 32 Grade 12 Chemistry
 Name: ___________________________________________

The following graph represents the equilibrium established for the chemical reaction,
Fe3+ + SCN- ↔ FeSCN2+
Concentration versus Time
8

6
Concentration (mol/L)

0 5 10 15 20 25 30 35

Time (s)

Fe3+ SCN- FeSCN2+

Use the graph for the next two questions.

5. At what time did this system first reach equilibrium?
a. 5 s
b. 10 s
c. 15 s
d. 20 s
6. What was the stress that occurred at 14 seconds?
a. the addition of Fe3+
b. the addition of SCN-
c. the removal of FeSCN2+
d. the addition of a catalyst
7. Which of the following does not describe a use of the Haber process?
a. the creation of nitrogen-containing explosives in WWI
b. the production of ammonia for household cleaners such as Windex
c. the creation of ammonia for industrial and household fertilizers
d. the distillation of petroleum products

Final Practise Examination Answer Key 3 of 32

 Acids and Bases (9 marks)

8. Which of the following statements does not support Arrhenius’ work with acids and
bases?
a. Arrhenius discovered that all acidic and basic solutions he tested were
electrolytes.
b. He determined that acids and bases must ionize or dissociate in water.
c. According to Arrhenius, an acid is defined as a hydrogen-containing compound
that ionizes to yield hydrogen ions (H+) dissolved in aqueous solution.
d. According to Arrhenius, an acid is defined as a compound that ionizes to yield
hydroxide ions (OH-) in aqueous solution.
9. Which of the following is considered a limitation of the Brønsted—Lowry definition of
acids and bases?
a. The Brønsted—Lowry theory expands the definition of an acid and base to a
proton donor or acceptor.
b. Even though acids and bases can occur without water, there is still a requirement
for the presence of a solvent.
c. The Brønsted—Lowry theory cannot explain why some substances without
hydroxides, like ammonia, can act as bases.
d. The Brønsted—Lowry theory does not explain how substances without protons
can act as acids.
10. Which of these descriptions is not a characteristic of the Lewis theory?
a. This theory limits the number of compounds called acids, since any compound
that has one or more valence shell orbitals cannot act as an acid.
b. The Lewis definition is so general that any reaction in which a pair of electrons is
transferred becomes an acid-base reaction.
c. Lewis acid-base reactions include many reactions that would not be included with
the Brønsted—Lowry definition.
d. The Lewis acid-base theory does not affect Brønsted—Lowry bases because any
Brønsted—Lowry base must have a pair of non-bonding electrons in order to
accept a proton.
11. At 25°C, the concentrations of the hydronium and hydroxide ions in water are equal at
a. 1.0 x 107 mol/L
b. 1.0 x 1014 mol/L
c. 1.0 x 10-7 mol/L
d. 1.0 x 10-14 mol/L

4 of 32 Grade 12 Chemistry
Name: ___________________________________________

12. Which statement regarding indicators is false?

a. They are strong, organic acids that change colour when the hydronium or
hydroxide ion concentration is changed.
b. Indicators change colour over a given pH range (the colour of each indicator can
be coordinated with the pH).
c. The colour of the indicator can be compared to a standard to determine the pH of
the solution.
d. Using an indicator is not as accurate as a pH metre.
13. All of these are examples of strong acids with the exception of
a. HClO4 (perchloric acid)
b. HI (hydroiodic acid)
c. HBr (hydrobromic acid)
d. H2CO3 (carbonic acid)
14. A carefully controlled neutralization reaction is also known as a/an
a. oxidation
b. reduction
c. titration
d. dissociation
15. On a titration curve, what indicates when neutralization of the unknown has
occurred?
a. The beginning of the horizontal region in the acid pH range.
b. The beginning of the horizontal region in the basic pH range.
c. The point halfway on the vertical region of the graph.
d. The beginning of the vertical region of the graph.
16. All of the following salts are BASIC in water with the exception of
a. sodium nitrate (NaNO3)
b. sodium acetate (NaC2H3O2)
c. potassium phosphate (K3PO4)
d. sodium sulfide (Na2S)

Final Practise Examination Answer Key 5 of 32

 Electrochemistry (7 marks)

17. Using the Activity Series, determine which species is the strongest oxidizing agent.
a. F- (aq)
b. F2 (g)
c. Li+ (aq)
d. Li (s)
18. Which type of cell converts chemical energy into electrical energy from a spontaneous
redox reaction?
a. an electrolytic cell
b. a half-cell
c. a wet cell
d. a voltaic cell
19. Which statement about an anode in a voltaic cell is false?
a. The anode is the electrode at which oxidation occurs.
b. The anode is the electrode at which reduction occurs.
c. Electrons are lost from this electrode.
d. It is normally labeled the negative electrode.
20. Which of the following components is not part of a voltaic cell?
a. a battery
b. two half-reaction cells
c. an external wire
d. the salt bridge
21. Which of the following statements related to cell potential is false?
a. The cell potential for a voltaic cell is a measure of the amount of voltage that can
be generated by driving an electric current through a wire.
b. One joule of energy is produced when one coulomb of electrical charge is
transported across a potential of one volt.
c. This difference in potential difference is an indication of how much energy is
available to move electrons from the anode to the cathode.
d. Electric charge can only flow between two electrodes when there is an equal
amount of electric potential between the two points.

6 of 32 Grade 12 Chemistry
Name: ___________________________________________

22. A short-hand notation is often used to represent the voltaic cell diagram. What does
the double vertical line represent?
a. It represents a phase boundary between the metal and the ion in solution.
b. It represents the salt bridge.
c. It represents the anode.
d. It represents the cathode.
23. Many ‘reference’ electrodes were tried before one was chosen as the standard to
which all other electrodes would be measured. Which half-cell was finally chosen as
the reference?
a. oxygen
b. copper
c. hydrogen
d. zinc

Final Practise Examination Answer Key 7 of 32

 Part B: Fill-in-the-Blank (20 marks total)

Use the answer sheet found at the end of this examination to answer the
fill-in-the-blank questions of this section. Write your answer in the space provided
that corresponds to the question. DO NOT write your answers directly on the
examination.
Using a term from the word bank provided below, complete each of the statements
that follow. Each blank is worth one mark; therefore, some questions have a total
value of two marks. There are MORE terms provided than you need, so read over the
list carefully and choose the terms you want to use.

acid energy mass action

acidic equal neutral
amphoteric equilibrium oxidation
anode equilibrium constant oxidizing
Arrhenius equivalence pH
base exothermix pOH
basic forward pressure
brine Galvani proton
cathode heterogeneous reducing
closed homogeneous reduction
closed system hydrolysis reversible
concentration hydronium salt bridge
constant increases spontaneous
Daniell indicator(s) standard reduction
Davy Kc potentials

dynamic Keq strong

electrolysis Lavoisier temperature

electrolyte Le Châtelier’s Principle titration

electrons Lewis Volta

electroplating Liebig

8 of 32 Grade 12 Chemistry
Name: ___________________________________________

Equilibrium (6 marks)
1. Equilibrium can only be established when all particles are kept in a sealed container
and certain conditions are kept ______________.
2. In a reversible reaction, when the rate at which products are formed equals the rate at
which reactants are formed, we say that the reaction has reached a ______________
equilibrium.
3. Equilibrium constants are specific for only one reaction at a particular ____________.
4. Le Châtelier’s Principle states that the equilibrium of a ______________
______________ will shift so as to relieve the stress, thereby restoring equilibrium.
5. If the forward reaction is endothermic, increasing the temperature ______________
the value of the equilibrium constant.
Acids and Bases (8 marks)
6. ______________ studied the conductivity of solutions and proposed that electrolytes
break up into charged particles in water.
7. When water acts as the base, accepting the proton, the result is the H3O+ ion called
the ______________ ion.
8. A conjugate base is what remains after a/an ______________ has donated its proton.
9. A neutral solution occurs when the hydronium ion concentration is ______________
to the hydroxide ion concentration.
10. In general, a strong acid will dissociate close to 100% and have a very large
______________.
11. An ______________ is defined as a compound that conducts an electric current when
it is in an aqueous state.
12. At equivalence, the ______________ ion and hydroxide ion concentrations are equal.
13. ______________ salts result from the reaction of a strong base with a weak acid.
Electrochemistry (6 marks)
14. A spontaneous reaction is one that occurs without any added ______________.
15. The stronger the oxidizing agent, the weaker the ______________ ability.
16. In an electrochemical cell, the electrode at which reduction occurs is called the
______________.
17. ______________ flow in the external circuit from the anode to the cathode of an
electrochemical cell.
18. All the half reactions on the Table of Standard Reduction Potentials with Values are
written as ______________.
19. ______________ is a process where an electric current is used to plate a metal onto
another surface.

Final Practise Examination Answer Key 9 of 32

 Part C: Short Answer (57 marks total)

Answer each of the following questions using the space provided. Pay attention to
the number of marks that each question is worth, as this may help you decide how
much information to provide for full marks. For questions that involve calculations,
show your work and check your final answer for the correct number of significant
figures as well as the appropriate unit.

Equilibrium (13 marks)

1. Explain how reaction rate and equilibrium are related concepts. Give an example to
illustrate this relationship. (2 marks)
When the rates of the forward and the reverse reactions are equal, equilibrium
has been established. (1 mark) For example, when a liquid is placed in a closed
container, equilibrium is established when the liquid vaporizes and condenses at
the same rate. (1 mark)

2. Write the equilibrium law (mass action expression) for the following reaction.
(2 marks)
2 SO2(g) + O2(g) ↔ 2 SO3(g)

[SO ]
2
3
K=
[SO ] [O ]
c 2
2 2

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Name: ___________________________________________

3. Explain what an equilibrium constant of 1 indicates about the reaction. (2 marks)

If K = 1, the ratio of [products] = [reactants] at equilibrium. (1 mark)
Neither reactants nor products are favoured at equilibrium. (1 mark)

4. For the reaction, CO(g) + 2 H2(g) ↔ CH3OH(g) + energy, predict the effect of the
following changes on the equilibrium concentration of CH3OH(g):
a. an increase in temperature (1 mark)
An increase in temperature decreases [CH3OH].

b. a decrease in pressure (1 mark)

A decrease in pressure decreases [CH3OH].

c. the addition of CO(g) (1 mark)

An addition of CO(g) increases [CH3OH].

Final Practise Examination Answer Key 11 of 32

 5. How might Le Châtelier’s Principle be useful in the chemical industry? For example,
how could you ensure a high yield in the production of ammonia? (2 marks)
Le Châtelier’s Principle could be used to maximize the amount of product formed
in chemical reactions. (1 mark) By maintaining high temperature and pressure, by
constantly adding reactants, and by immediately removing products, the products
continue to be favoured. (1 mark)

6. Several types of stress can disrupt chemical equilibrium. Name two such stresses.
(2 marks)
Any two of the following for 1 mark each:
QQ changes in the concentration of reactants or products
QQ temperature changes
QQ pressure changes
QQ volume changes
QQ the addition of a catalyst

12 of 32 Grade 12 Chemistry
Name: ___________________________________________

Acids and Bases (23 marks)

7. Identify the acid, base, conjugate acid, and conjugate base in the following reversible
reaction. (2 marks)
HF(aq) + HSO3-(aq) ↔ F-(aq) + H2SO3(aq)
acid: HF(aq)
base: HSO3-(aq)
conjugate acid: H2SO3(aq)
conjugate base: F-(aq)

8. Using the ionization of water equation given, predict the effect of dissolving a base on
hydronium and hydroxide ion concentrations by using Le Châtelier’s Principle.
(3 marks)
H2O(l) + H2O(l) ↔ H3O+(aq) + OH-(aq)
When a base is dissolved in water, the hydroxide ion concentration increases.
(1 mark) According to Le Châtelier’s Principle, the equilibrium shifts left to use
up some of the added hydroxide and maintain Kw at 1.0 x 10-14. (1 mark) Since
equilibrium shifts left, the hydronium ion concentration is reduced. (1 mark) As
such, adding a base to water increases the hydroxide ion concentration and reduces
the hydronium ion concentration.

Final Practise Examination Answer Key 13 of 32

 9. What is the [H3O+] in a solution with [OH–] of 5.67×10–3? (2 marks)

KW 1.0×10 -14
[H3O+ ] = -
= -3
= 1.76×10 -12 mol/L H3O+
[OH ] 5.67×10

10. In a neutral solution, the [OH-] = 1 x 10-7 mol/L. Calculate the pOH of a neutral
solution, showing all the steps of your work. (2 marks)
pOH = –log [OH-] (1 mark)
pOH = –(log OH-) = (–log 10-7) = 7.0 (1 mark)

11. Are all strong acids also strong electrolytes? Explain your answer. (2 marks)
Yes, all strong acids are also by definition strong electrolytes. In water, strong
electrolytes completely dissociate into ions, and become good conductors of
electricity.

14 of 32 Grade 12 Chemistry
Name: ___________________________________________

12. Calculate the percent dissociation of a 0.20 mol/L solution of the weak acid, HNO2, if
the pH of the solution is 4.20. The dissociation reaction has been provided. Show all
the steps of your work. (5 marks)
HNO2(aq) + H2O(l) → H3O+(aq) + NO2-(aq)
[H3O+] = 10-pH = 10- 4.20 = 6.3 x 10-5 mol/L (2 marks)

[H3O+] 6.3 × 10-5 mol/L

% dissociation = ——— —— × 100 = ——————————— × 100 (2 marks)
[HNO2] 0.2 mol/L

% dissociation = 0.0315 % (1 mark)

Final Practise Examination Answer Key 15 of 32

 13. Calculate the concentration of the acid, if 25.0 mL of H3PO4 is required to neutralize
29.0 mL of 0.830 mol/L KOH. Complete the neutralization reaction and then show all
the steps of your work. (7 marks)
H3PO4(aq) + 3 KOH(aq) →
H3PO4(aq) + 3 KOH(aq) → K3PO4(aq) + 3 H2O(l) (1 mark)
moles of KOH = C x V = (0.830 mol/L)(0.0290 L) = 0.02407 moles KOH (2 marks)
 1 mole H3 PO4 
moles H3 PO4 = 0.02407 moles KOH  = 0.008023 moles H3 PO4 (2 marks)
 3 mole KOH 

mol 0.008023 moles

Concentration = = = 0.321 mol/L (2 marks)
V 0.0250 L

The student may have constructed the following cvn table:

3 KOH H3PO4

C (mol/L) 0.830 mol/L 0.321 mol/L

V (L) 0.029 L 0.025 L

n (mol) 0.02407 mol 0.008023 mol

16 of 32 Grade 12 Chemistry
Name: ___________________________________________

Electrochemistry (21 marks)

14. Explain how the activity series can be used to predict if a reaction will be
spontaneous or not. (2 marks)
The spontaneity rule states that any species on the left side of the activity series
will react spontaneously and oxidize any species on the right side which is above
it. In other words, “spontaneous reactions are up-right.”

15. Two half-cells are connected under standard conditions to make an

electrochemical cell. The two half-cells are a copper-copper(I) ion (Cu/Cu+) and
an aluminum-aluminum ion (Al/Al3+). Using your Table of Standard Reduction
Potentials with Values, complete the following:
a. Write the two half reactions that will occur within this electrochemical cell.
(2 marks)
1 mark for each half reaction, even if is not labelled.
Cu+ + e– → Cu(s)
Al(s) → Al3+(aq) + 3e–

b. Identify which reaction is the anode and which is the cathode .

(0.5 mark x 2 = 1 mark)
Cathode: Cu+ + e– → Cu(s) (0.5 mark)
Anode: Al(s) → Al3+(aq) + 3e– (0.5 mark)

Final Practise Examination Answer Key 17 of 32

 c. Calculate the overall voltage of the cell. (2 marks)
+0.52V + 1.66V = +2.18V

d. Write the overall balanced redox reaction for the electrochemical cell. (3 marks)
(x3) Cu+ + e- → Cu(s)
Al(s) → Al3+(aq) + 3e-
__________________________________________
Al(s) + 3 Cu+(aq) → Al3+(aq) + 3 Cu(s)

18 of 32 Grade 12 Chemistry
Name: ___________________________________________

e. Draw a labelled electrochemical cell indicating

—— the electrodes (anode, cathode, positive, and negative) (2 marks)
—— the types of ions (2 marks)
—— the direction of electron flow (0.5 mark)
—— the electrode half reactions (2 marks)
—— the salt bridge (0.5 mark)

0

e- e-
- +

- * +
NO3- K+

Anode Cathode
- post + post
Al electrode Cu electrode

Al (NO3 )3 CuNO 3
Al Cu

Cu+
Al 3+

NO3- NO3-

Al (s) → Al3+(aq) + 3e - Cu +(aq) + e - → Cu (s)

Anode Oxidation Cathode Reduction

*Note: Any combination of cation and anion

is acceptable in the salt bridge.

Final Practise Examination Answer Key 19 of 32

 16. Calculate how long it would take an aqueous gold (III) chloride cell to plate 2.5 g of
gold on a bracelet using a current of 2.5 A. The half reaction has been provided for
you. (4 marks)
Half reaction: Au3+(aq) + 3e– → Au(s)

m 2.50 g
moles Au = = = 0.01269 moles Au
M 197.0 g/mol
 3 moles e - 
moles e - = 0.01269 moles Au   = 0.03807 moles e -
 1 mole Zn 

(96500 C/mol ) ne (96500 C/mol ) (0.03807 moles e- )

-
t= = = 1469.5 s
I 2.5 A

It would require about 1470 s (24.5 minutes) to plate 2.5 g of gold.

20 of 32 Grade 12 Chemistry

Group
1 18

1 Atomic 19 2
H Number He
K Symbol
Hydrogen Helium
1.0 2 Name Potassium 4.0
1 13 14 15 16 17 1
39.1 Relative
3 4 Atomic Mass 5 6 7 8 9 10
Li Be B C N O F Ne
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
6.9 9.0 10.8 12.0 14.0 16.0 19.0 20.2
2 2
11 12 13 14 15 16 17 18
Na Mg Al Si P S Cl Ar
Sodium Magnesium Aluminum Silicon Phosphorus Sulphur Chlorine Argon
23.0 24.3 3 4 5 6 7 8 9 10 11 12 27.0 28.1 31.0 32.1 35.5 39.9
3 3
19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krpton
39.1 40.1 45.0 47.9 50.9 52.0 54.9 55.8 58.9 58.7 63.5 65.4 69.7 72.6 74.9 79.0 79.9 83.8
4 4
37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
85.5 87.6 88.9 91.2 92.9 96.0 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3
5 5
55 56 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86
57–71
Cs Ba Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Lanthanide
Cesium Barium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury Thallium Lead Bismuth Polonium Astatine Radon
Series
132.9 137.3 178.5 180.9 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222)
6 6
87 88 104 105 106 107 108 109 110 111 112 113 114 115 116 118
89–103
Name: ___________________________________________

Fr Ra Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uuo

Actinide
Francium Radium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadium Roentgenium Copernicium Ununtrium Ununquadium Ununpentium Ununhexium Ununoctium
Series
(223) (226) (261) (268) (271) (272) (270) (276) (281) (280) (285) (284) (289) (288) (293) (294)
7 7
Periodic Table of Elements

57 58 59 60 61 62 63 64 65 66 67 68 69 70 71
Lanthanide Series La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium
Inner 138.9 140.1 140.9 144.2 (145) 150.4 152.0 157.2 158.9 162.5 164.9 167.3 168.9 173.0 174.9
Transition
Elements
89 90 91 92 93 94 95 96 97 98 99 100 101 102 103
Actinide Series Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

Final Practise Examination Answer Key

Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium
(227) 232.0 231.0 238.0 (237) (244) (243) (247) (247) (251) (252) (257) (258) (259) (262)

21 of 32
 Alphabetical Listing of the Elements and Their Atomic Masses

Element Atomic Mass Element Atomic Mass Element Atomic Mass

Actinium (227) Gold 197.0 Praseodymium 140.9
Aluminum 27.0 Hafnium 178.5 Promethium (145)
Americium (243) Hassium (265) Protactinum (231)
Antimony 121.7 Helium 4.0 Radium (226)
Argon 39.9 Holmium 164.9 Radon (222)
Arsenic 74.9 Hydrogen 1.0 Rhenium 186.2
Astatine (210) Indium 114.8 Rhodium 102.9
Barium 137.3 Iodine 126.9 Rubidium 85.5
Berkelium (247) Irdium 192.2 Ruthenium 101.1
Beryllium 9.0 Iron 55.8 Rutherfordium (261)
Bismuth 209.0 Krypton 83.8 Samarium 150.4
Bohrium (264) Lanthanum 138.9 Scandium 45.0
Boron 10.8 Lawrencium (257) Seaborgium (263)
Bromine 79.9 Lead 207.2 Selenium 79.0
Cadmium 112.4 Lithium 6.9 Silicon 28.1
Calcium 40.1 Lutetium 175.0 Silver 107.9
Californium (251) Magnesium 24.3 Sodium 23.0
Carbon 12.0 Manganese 54.9 Strontium 87.6
Cerium 140.1 Meitnerium (266) Sulfur 32.1
Cesium 132.9 Mendelevium (256) Tantalum 180.9
Chlorine 35.5 Mercury 200.6 Technetium (98)
Chromium 52.0 Molybdenum 95.9 Tellurium 127.6
Cobalt 58.9 Neodymium 144.2 Terbium 158.9
Copernicium (277) Neon 20.2 Thallium 204.4
Copper 63.5 Neptunium (237) Thorium 232.0
Curium (247) Nickel 58.7 Thulium 168.9
Dubnium (262) Niobium 92.9 Tin 118.7
Dysprosium 162.5 Nitrogen 14.0 Titanium 47.9
Einstienium (254) Nobelium (259) Tungsten 183.8
Erbium 167.3 Osmium 190.2 Uranium 238.0
Europium 152.0 Oxygen 16.0 Vanadium 50.9
Fermium (257) Palladium 106.4 Xenon 131.3
Fluorine 19.0 Phosphorus 31.0 Ytterbium 173.0
Francium (223) Platinum 195.1 Yttrium 88.9
Gadolinium 157.2 Plutonium (244) Zinc 65.4
Gallium 69.7 Polonium (209) Zirconium 91.2
Germanium 72.6 Potassium 39.1

22 of 32 Grade 12 Chemistry

Group
1 18

1 2
H He
2.20 —
1 2 13 14 15 16 17 1

3 4 5 6 7 8 9 10
Li Be B C N O F Ne
0.97 1.47 2.01 2.50 3.07 3.50 4.10 —
2 2

11 12 13 14 15 16 17 18
Na Mg Al Si P S Cl Ar
1.01 1.23 1.47 1.74 2.06 2.44 2.83 —
3 3 4 5 6 7 8 9 10 11 12 3

19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
0.91 1.04 1.20 1.32 1.45 1.56 1.60 1.64 1.70 1.75 1.75 1.66 1.82 2.02 2.20 2.48 2.74 —
4 4

37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
0.89 0.99 1.11 1.22 1.23 1.30 1.36 1.42 1.45 1.35 1.42 1.46 1.49 1.72 1.82 2.01 2.21 —
5 5

55 56 57–71 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86
Cs Ba Lanthanide Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
0.86 0.97 Series 1.23 1.33 1.40 1.46 1.52 1.55 1.44 1.42 1.44 1.44 1.55 1.67 1.76 1.90 —
6 6

87 88 104 105 106 107 108 109 110 111 112 113 114 115 116 118
Name: ___________________________________________

89–103
Fr Ra Actinide Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uuo
0.86 0.97 Series — — — — — — — — — — — — — —
7 7
Table of Electronegativities

57 58 59 60 61 62 63 64 65 66 67 68 69 70 71
Lanthanide Series La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Inner 1.08 1.08 1.07 1.07 1.07 1.07 1.01 1.11 1.10 1.10 1.10 1.11 1.11 1.06 1.14
Transition
Elements
89 90 91 92 93 94 95 96 97 98 99 100 101 102 103
Actinide Series

Final Practise Examination Answer Key

Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
1.00 1.11 1.14 1.30 1.29 1.25 — — — — — — — — —

23 of 32
 Relative Strengths of Acids Table

Acid Reaction Ka
Perchloric acid HClO4 + H2O → H3O+ + ClO4– very large
Hydriodic acid Hl + H2O → H3O + I
+ –
very large
Hydrobromic acid HBr + H2O → H3O+ + Br – very large
Hydrochloric acid HCl + H2O → H3O+ + Cl– very large
Nitric acid HNO3 + H2O → H3O+ + NO3– very large
Sulfuric acid H2SO4 + H2O → H3O + HSO4
+ –
very large
Oxalic acid H2C2O4 + H2O → H3O + HC2O4
+ –
5.4 x 10–2
Sulfurous acid H2SO3 + H2O → H3O+ + HSO3– 1.7 x 10–2
Hydrogen sulfate ion HSO4– + H2O → H3O+ + SO42– 1.3 x 10–2
Phosphoric acid H3PO4 + H2O → H3O+ + H2PO4– 7.1 x 10–3
Ferric ion Fe(H2O)63+ + H2O → H3O+ + Fe(H2O)5(OH)2+ 6.0 x 10–3
Hydrogen telluride H2Te + H2O → H3O+ + HTe– 2.3 x 10–3
Hydrofluoric acid HF + H2O → H3O+ + F– 6.7 x 10–4
Nitrous acid HNO2 + H2O → H3O+ + NO2– 5.1 x 10–4
Hydrogen selenide H2Se + H2O → H3O+ + HSe– 1.7 x 10–4
Chromic ion Cr(H2O)63+ + H2O → H3O+ + Cr(H2O)5(OH)2+ 1.5 x 10–4
Benzoic acid C6H5COOH + H2O → H3O+ + C6H5COO– 6.6 x 10–5
Hydrogen oxalate ion HC2O4– + H2O → H3O+ + C2O42– 5.4 x 10–5
Acetic acid HC2H3O2 + H2O → H3O+ + C2H3O2– 1.8 x 10–5
Aluminum ion Al(H2O)63+ + H2O → H3O+ + Al(H2O)5(OH)2+ 1.4 x 10–5
Carbonic acid H2CO3 + H2O → H3O+ + HCO3– 4.4 x 10–7
Hydrogen sulfide H2S + H2O → H3O+ + HS– 1.0 x 10–7
Dihydrogen phosphate ion H2PO4– + H2O → H3O+ + HPO42– 6.3 x 10–8
Hydrogen sulfite ion HSO3– + H2O → H3O+ + SO32– 6.2 x 10–8
Ammonium ion NH4+ + H2O → H3O+ + NH3 5.7 x 10–10
Hydrogen carbonate ion HCO3– + H2O → H3O+ + CO32– 4.7 x 10–11
Hydrogen telluride ion HTe– + H2O → H3O+ + Te2– 1.0 x 10–11
Hydrogen peroxide H2O2 + H2O → H3O+ + HO2– 2.4 x 10–12
Monohydrogen phosphate HPO42– + H2O → H3O+ + PO43– 4.4 x 10–13
Hydrogen sulfide ion HS– + H2O → H3O+ + S2– 1.2 x 10–15
Water H2O + H2O → H3O+ + OH– 1.8 x 10–16
Hydroxide ion OH– + H2O → H3O+ + O2– < 10–36
Ammonia NH3 + H2O → H3O+ + NH2– very small

24 of 32 Grade 12 Chemistry
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Solubility Chart

Negative Ions Positive Ions Solubility

essentially all alkali ions (Li+, Na+, K+, Rb+, Cs+) soluble

essentially all hydrogen ion H+(aq) soluble

essentially all ammonium ion (NH4+) soluble

nitrate, NO3– essentially all soluble

acetate, CH3COO– essentially all (except Ag+) soluble

Ag+, Pb2+, Hg22+, Cu+, Tl+ low solubility

chloride, Cl–
bromide, Br–
iodide, I–
all others soluble

Ca2+, Sr2+, Ba2+, Pb2+, Ra2+ low solubility

sulfate, SO42–
all others soluble

alkali ions, H+(aq), NH4+, Be2+, Mg2+, Ca2+,

soluble
Sr2+, Ba2+, Ra2+
sulfide, S2–

all others low solubility

alkali ions, H+(aq), NH4+, Sr2+, Ba2+, Ra2+, Tl+ soluble

hydroxide, OH –

all others low solubility

alkali ions, H+(aq), NH4+ soluble

phosphate, PO43–
carbonate, CO32–
sulfite, SO32–
all others low solubility

Ba2+, Sr2+, Pb2+, Ag+ low solubility

chromate, CrO42–
all others soluble

Final Practise Examination Answer Key 25 of 32

 Table of Standard Reduction Potentials with Values

Oxidized species ↔ Reduced Species E°/V

Li+(aq) + e– ↔ Li(s) –3.04

K+(aq) + e– ↔ K(s) –2.93

Ca2+(aq) + 2e– ↔ Ca(s) –2.87

Na+(aq) + e– ↔ Na(s) –2.71

Mg2+(aq) + 2e– ↔ Mg(s) –2.37

Al3+(aq) + 3e– ↔ Al(s) –1.66

Mn2+(aq) + 2e– ↔ Mn(s) –1.19

H2O(l) + e– ↔ ½H2(g) + OH–(aq) –0.83

Zn2+(aq) + 2e– ↔ Zn(s) –0.76

Fe2+(aq) + 2e– ↔ Fe(s) –0.45

Ni2+(aq) + 2e– ↔ Ni(s) –0.26

Sn2+(aq) + 2e– ↔ Sn(s) –0.14

Pb2+(aq) + 2e– ↔ Pb(s) –0.13

H+(aq) + e– ↔ ½H2(g) 0.00

Cu2+(aq) + e– ↔ Cu+(aq) +0.15

SO42–(aq) + 4 H+(aq) + 2e– ↔ H2SO3(aq) + H2O(l) +0.17

Cu2+(aq) + 2e– ↔ Cu(s) +0.34

½O2(g) + H2O(l) + 2e– 2OH–(aq) +0.40

Cu+(aq) + e– Cu(s) +0.52

½I2(s) + e– ↔ I–(aq) +0.54

Fe3+(aq) + e– ↔ Fe2+(aq) +0.77

Ag+(aq) + e– ↔ Ag(s) +0.80

½Br2(l) + e– ↔ Br – (aq) +1.07

½O2(g) + 2H+(aq) + 2e– ↔ H2O(l) +1.23

Cr2O72–(aq) + 14H+(aq) + 6e– ↔ 2Cr3+(aq) + 7H2O(l) +1.33

½Cl2(g) + e– ↔ Cl– (aq) +1.36

MnO4–(aq) + 8H+(aq) + 5e– ↔ Mn2+(aq)+ 4H2O(l) +1.51

½F2(g) + e– ↔ F– (aq) +2.87

26 of 32 Grade 12 Chemistry
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Names, Formulas, and Charges of Common Ions

Positive Ions (Cations)

Name Symbol Name Symbol
aluminum Al3+ magnesium Mg2+
ammonium NH4+ manganese(II) Mn2+
barium Ba2+ manganese(IV) Mn4+
cadmium Cd2+ mercury(I) Hg22+
calcium Ca2+ mercury(II) Hg2+
chromium(II) Cr2+ nickel(II) Ni2+
chromium(III) Cr3+ nickel(III) Ni3+
copper(I) Cu+ potassium K+
copper(II) Cu2+ silver Ag+
hydrogen H+ sodium Na+
iron(II) Fe2+ strontium Sr2+
iron(III) Fe3+ tin(II) Sn2+
lead(II) Pb2+ tin(IV) Sn4+
lead(IV) Pb4+ zinc Zn2+
lithium Li+
continued

Final Practise Examination Answer Key 27 of 32

 Names, Formulas, and Charges of Common Ions (continued)

Negative Ions (Anions)

Name Symbol Name Symbol
acetate C2H3O2—(CH3COO—) nitrate NO3—
azide N3 — nitride N3—
bromide Br— nitrite NO2—
bromate BrO3— oxalate C2O42—
carbonate CO32— hydrogen oxalate HC2O4—
hydride H— oxide O2—
hydrogen carbonate
HCO3— perchlorate ClO4—
or bicarbonate
chlorate ClO3— permanganate MnO4—
chloride Cl— phosphate PO43—
chlorite ClO2— monohydrogen phosphate HPO42—
chromate CrO42— dihydrogen phosphate H2PO4—
citrate C6H5O73— silicate SiO32—
cyanide CN— sulfate SO42—
dichromate Cr2O72— hydrogen sulfate HSO4—
fluoride F— sulfide S2—
hydroxide OH— hydrogen sulfide HS—
hypochlorite ClO— sulfite SO32—
iodide I— hydrogen sulfite HSO3—
iodate IO3— thiocyanate SCN—

28 of 32 Grade 12 Chemistry
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Common Ions

Cations (Positive Ions)

1+ charge 2+ charge 3+ charge
NH4+ Ammonium Ba2+ Barium Al3+ Aluminum
Cs+ Cesium Be2+ Beryllium Cr3+ Chromium(III)
Cu+ Copper(I) Cd2+ Cadmium Co3+ Cobalt(III)
Au+ Gold(I) Ca2+ Calcium Ga3+ Gallium
H+ Hydrogen Cr2+ Chromium(II) Au3+ Gold(III)
Li+ Lithium Co2+ Cobalt(II) Fe3+ Iron(III)
K+ Potassium Cu2+ Copper(II) Mn3+ Manganese
Rb+ Rubidium Fe2+ Iron(II) Ni3+ Nickel(III)
Ag+ Silver Pb2+ Lead(II)
Na+ Sodium Mg2+ Magnesium 4+ charge
Mn2+ Manganese(II) Pb4+ Lead(IV)
Hg22+ Mercury(I) Mn4+ Manganese(IV)
Hg2+ Mercury(II) Sn4+ Tin(IV)
Ni2+ Nickel(II)
Sr2+ Strontium
Sn2+ Tin(II)
Zn2+ Zinc
(continued)

Final Practise Examination Answer Key 29 of 32

 Common Ions (continued)

Anions (Negative Ions)

1— charge 1— charge 2— charge

CH3COO— Acetate (or HS— Hydrogen CO32— Carbonate

(C2H3O2—) ethanoate) sulfide CrO42— Chromate
BrO3— Bromate OH— Hydroxide Cr2O72— Dichromate
Br— Bromide IO3— Iodate O2— Oxide
ClO3— Chlorate I— Iodide O22— Peroxide
Cl— Chloride NO3— Nitrate SO42— Sulfate
ClO2— Chlorite NO2— Nitrite S2— Sulfide
CN— Cyanide ClO4— Perchlorate SO32— Sulfite
F— Fluoride IO4— Periodate S2O32— Thiosulfate
H— Hydride MnO4— Permanganate
HCO3— Hydrogen SCN— Thiocynate 3— charge
carbonate (or
bicarbonate) N3— Nitride
ClO— Hypochlorite PO43— Phosphate
HSO4— Hydrogen P3— Phosphide
sulfate PO33— Phosphite

30 of 32 Grade 12 Chemistry
Name: ___________________________________________

Final Practice Examination Answer Sheet

Part A: Multiple Choice (23 marks)

For each multiple-choice question, shade in the circle that corresponds to your answer.
DO NOT circle your answers directly on the examination.
Example: A B C D

1. A B C D 9. A B C D 17. A B C D

2. A B C D 10. A B C D 18. A B C D

3. A B C D 11. A B C D 19. A B C D

4. A B C D 12. A B C D 20. A B C D

5. A B C D 13. A B C D 21. A B C D

6. A B C D 14. A B C D 22. A B C D

7. A B C D 15. A B B D 23. A B C D

8. A B C C 16. A B C D

Part B: Fill-in-the-Blank (20 marks)

For each fill-in-the-blank question, write your answer in the space provided that
corresponds to the question. DO NOT write your answers directly on the examination.
Equilibrium (6 marks)

1. constant
2. dynamic
3. temperature
4. closed
system
5. increases

continued

Final Practise Examination Answer Key 31 of 32

 Acids and Bases (8 marks)

6. Arrhenius
7. hydronium
8. acid
9. equal
10. Keq
11. electrolyte
12. hydronium
13. basic

Electrochemistry (6 marks)

14. energy
15. reducing
16. cathode
17. Electrons
18. reductions
19. Electroplating

32 of 32 Grade 12 Chemistry