R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 1 / 69
Chemistry 12
Provincial Exam Workbook
Unit 04: Acid Base Equilibria
Multiple Choice Questions
1.
Calculate the volume of 0.300 M HNO3 needed to completely neutralize 25.0 mL of
0.250 M Sr  OH2 .
A. 10.4 mL
B. 15.0 mL
C. 20.8 mL
D. 41.7 mL
2.
A.
B.
C.
D.
3.
Equal moles of which of the following chemicals could be used to make a basic buffer
solution?
HF and NaOH
HCl and NaCl
KBr and NaNO3
NH3 and NH4 Cl
Which reaction occurs when calcium oxide is added to water?
A. 2CaO(s)  Ca2O2 (aq)
B. 2CaO(s)  2Ca(aq)  O2 (aq)
C. CaO(s)  H2O(l)  Ca  OH2 (aq)
D. CaO(s)  H2O(l)  CaOH(aq)  O2 (aq)
4.
In which of the following is HSO3  acting as a Brönsted-Lowry acid?
A. HSO3  H2O  H2SO3  OH
B. NH3  HSO3  NH4  SO32
C. HSO3  HPO42  H2SO3  PO43
D. H2 C2O4  HSO3   HC2 O4   H2 SO3
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5.
What is the conjugate base of H2PO4  ?
A. OH
B. PO43
C. HPO42
D. H3PO4
6.
Which of the following is correct if the four solutions listed are compared to one
another?
7.
Which of the following is the strongest acid that can exist in an aqueous solution?
A. O2
B. NH2 
C. H3 O
D. HClO4
8.
Which of the following household products could have a pH = 12.0?
A. soda pop
B. tap water
C. lemon juice
D. oven cleaner
9.
What is the pH of a 0.050 M KOH solution?
A. 0.30
B. 1.30
C. 12.70
D. 13.70
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 3 / 69
10.
What is the value of K b for H2PO4  ?
A. 1.3x1012
B. 6.2x108
C. 1.6x107
D. 7.5x103
11.
Which of the following describes the net ionic reaction for the hydrolysis of NH4 Cl(s) ?
A. NH4 (aq)  Cl (aq)  NH4Cl(s)
B. NH4 Cl(s)  NH4  (aq)  Cl (aq)
C. Cl (aq)  H2 O(l)  HCl(aq)  OH (aq)
D. NH4 (aq)  H2O(l)  NH3 (aq)  H3O (aq)
12.
Which of the following salts will produce a solution with the highest pH?
A. 1.0 M
B. 1.0 M
C. 1.0 M
D. 1.0 M
13.
NaNO3
NaHSO4
NaHCO3
NaH2PO4
What is true about the transition point of all indicators described by the following
equilibrium…
HIn  H2O  H3 O  In
A. pH  K a
B. HIn  In 
C. H3 O   1.0x107M
D. moles of H3 O equals moles of OH
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 4 / 69
14.
15.
A chemical indicator has a Ka  4.0x106 . What is the pH at the transition point and
the identity of the indicator?
A 20.0 mL sample of HCl is titrated with 25.0 mL of 0.20 M Sr  OH2 . What is the
concentration of the acid?
A. 0.13 M
B. 0.20 M
C. 0.25 M
D. 0.50 M
16.
Which of the following is the net ionic equation for the neutralization of HNO3 (aq) with
Sr  OH2 (aq) ?
A. H (aq)  OH (aq)  H2O(l)
B. Sr 2 (aq)  2NO3 (aq)  Sr NO3 2 (aq)  2H2 O(l)
C. 2HNO3 (aq)  Sr  OH2 (aq)  Sr NO3 2 (aq)  2H2O(l)
D. 2H (aq)  2NO3 (aq)  Sr 2 (aq)  2OH (aq)  Sr 2 (aq)  2NO3 (aq)  2H2O(l)
17.
When a strong acid is titrated with a strong base, what will the pH value be at the
equivalence point?
A. 0.0
B. 6.8
C. 7.0
D. 8.6
18.
A.
B.
C.
D.
Which of the following acids could not be present in a buffer solution?
HF
HNO2
H2 SO3
HClO4
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 5 / 69
19.
Which net ionic equation best describes the reaction between NaOH and H2 S ?
A. OH (aq)  H (aq)  H2O(l)
B. 2OH (aq)  H2S(aq)  2H2O(l)  S2 (aq)
C. 2NaOH(aq)  H2S(aq)  2H2O(l)  Na2S(aq)
D. 2Na (aq)  2OH (aq)  2H (aq)  S2 (aq)  2H2 O(l)  2Na (aq)  S2 (aq)
20.
Which of the following is a general characteristic of Arrhenius acids?
A. They produce H in solution
B. They accept an H from water
C. They turn bromthymol blue a blue colour
D. They react with H3 O ions to produce H2
21.
Identify a conjugate pair from the equilibrium provided…
PO43  HCO3  HPO42  CO32
A. CO32 and PO43
B. PO43 and HCO3 
C. PO43- and HPO42D. HCO3  and HPO42
22.
Which of the following best describes a weak acid?
A. Its 0.10 M solution will have pH of 1.00
B. It may be very soluble, but only partly ionized
C. It must be very soluble and completely ionized
D. It must be of low solubility and completely ionized
23.
What is the pH at the transition point for an indicator with a K a of 2.5x104 ?
A. 2.5x104
B. 3.60
C. 7.00
D. 10.40
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24.
What volume of 0.100 M NaOH is required to completely neutralize 15.00 mL of
0.100 M H3PO4 ?
A. 5.00 mL
B. 15.0 mL
C. 30.0 mL
D. 45.0 mL
25.
What is the pH of the solution formed when 0.060 moles NaOH is added to 1.00 L of
0.050 M HCl ?
A. 2.00
B. 7.00
C. 12.00
D. 12.78
26.
Which of the following graphs describes the relationship between the pH of a buffer
and the volume of NaOH added to the buffer?
27.
A gas which is produced by internal combustion engines and contributes to the
formation of acid rain is…
A.
B.
C.
D.
H2
O3
CH4
NO2
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 7 / 69
28.
Consider the following reaction…
H3BO3 (aq)  HS (aq)  H2BO3  (aq)  H2S(aq)
The order of Brönsted-Lowry acids and bases in this equation is…
A. acid, base, base, acid
B. acid, base, acid, base
C. base, acid, acid, base
D. base, acid, base, acid
29.
The conjugate base of an acid is produced by…
A. adding a proton to the acid
B. adding an electron to the acid
C. removing a proton from the acid
D. removing an electron from the acid
30.
Which of the following represents the predominant reaction between HCO3  and
water?
A. 2HCO3  H2 O  2CO2
B. HCO3   H2 O  H2 CO3  OH
C. HCO3   H2 O  H3 O  CO32
D. 2HCO3  H2O  H3 O  CO32  OH  CO2
31.
Water acts as an acid when it reacts with which of the following?
A. I and IV only
B. II and III only
C. I, II and IV only
D. II, III and IV only
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32.
A 1.0x104M solution has a pH of 10.00. The solute is a…
A. weak acid
B. weak base
C. strong acid
D. strong base
33.
Consider the following Brönsted-Lowry equilibrium system…
HSO3  H2PO4   SO32  H3PO4
What are the two Brönsted-Lowry bases in the equilibrium above?
A. HSO3  and SO3 2
B. H2PO4  and SO32
C. HSO3  and H3PO4
D. H2PO4  and H3PO4
34.
Which of the following will have the lowest electrical conductivity?
A. 1.00M HCl
B. 1.00M HNO3
C. 1.00M H2 SO4
D. 1.00M H3PO4
35.
Which of the following represents the ionization of water?
A. H2 O  H2  0.5O2
B. 2H2 O  H3 O  OH
C. 2H2O  O2  2H2O2
D. H2 O  0.5O2  2H  2e
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36.
Which of the following are amphiprotic?
A. I and II only.
B. I and III only.
C. II and III only.
D. I, II and III.
37.
Consider the following equilibrium:
2H2O(l)  energy  H3 O (aq)  OH (aq)
The temperature is increased and a new equilibrium is established. The new equilibrium can
be described by…
A. pH  pOH and K W  1.0x1014
B. pH  pOH and K W  1.0x1014
C. pH  pOH and K W  1.0x1014
D. pH  pOH and K W  1.0x1014
38.
What is the H3 O  at the equivalence point for the titration between HBr and KOH ?
A. 1.0x109M
B. 1.0x107M
C. 1.0x105M
D. 0.0 M
39.
A.
B.
C.
D.
Which of the following would form a buffer solution when equal moles are mixed
together?
HCl and NaCl
HCN and NaCN
KNO3 and KOH
Na2SO4 and NaOH
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 10 / 69
40.
Which of the following oxides dissolves to form a solution with a pH greater than 7?
A. SO2
B. CO2
C. N2 O
D. K 2 O
41.
The pH of acid rain could be…
A. 5.0
B. 7.0
C. 9.0
D. 11.0
42.
A hydronium ion has the formula…
A. H2 
B. OH
C. H2 O
D. H3 O
43.
The conjugate acid of C6H5NH2 is…
A. C6H5NH
B. C6H5NH3
C. C6H5NH2
D. C6H5NH3 
44.
Which of the following is a property of 1.0 M HCl but not a property of 1.0 M
CH3 COOH ?
A. turns litmus red
B. ionizes completely
C. has a pH less then 7.0
D. produces H3 O in solution
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45.
In a 1.0 M HF solution, the concentration of HF , F , and OH , from highest to lowest
is…
A. HF  F   OH 
B. F   HF  OH 
C. OH   HF  F 
D. OH   F   HF
46.
In which of the following reactions is water behaving as a Brönsted-Lowry acid?
A. 2H2 O  2H2  O2
B. HCl  H2 O  H3 O  Cl
C. NH3  H2O  NH4   OH
D. NH4  H2O  H3 O  NH3
47.
A.
B.
C.
D.
What is the OH  of a solution with H3 O   9.3x102M ?
9.3x1016M
8.6x1013M
1.1x1013M
9.3x102M
48.
The pH of 0.10 M HNO3 is…
A. 0.79
B. 1.00
C. 1.26
D. 13.00
49.
What term is used to describe the point at which a chemical indicator changes
colour?
A. titration point
B. transition point
C. equivalence point
D. stoichiometric point
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50.
When 25.0 mL samples of the strong acid H2 SO4 were titrated with 0.25 M NaOH
the following results were obtained…
What is the concentration of the H2 SO4 sample?
A. 0.20 M
B. 0.21 M
C. 0.40 M
D. 0.42 M
51.
Which of the following equations describes the predominant reaction that occurs at
the equivalence point of a titration between CH3 COOH(aq) and NaOH(aq) ?
A. H (aq)  OH (aq)  H2O(l)
B. CH3 COO (aq)  H2O(l)  CH3 COOH(aq)  OH (aq)
C. CH3 COOH(aq)  NaOH(aq)  NaCH3 COO(aq)  H2O(l)
D. H (aq)  CH3 COO (aq)  Na (aq)  OH (aq)  Na (aq)  CH3COO (aq)  H2O(l)
52.
Why is the solution below not considered to be a true buffer solution?
HCN(aq)
1.0M
A. excessive HCN
B. excessive H3 O 
C. insufficient CN 
D. insufficient HCN
 H2 O(l) 
H3 O (aq)

CN (aq)
2.2x105M 2.2x107M
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53.
What could be added to 1.0 L of the solution below in order for it to behave as a true
buffer?
HCN(aq)
1.0M
A. 1.0 mol
B. 1.0 mol
C. 1.0 mol
D. 1.0 mol
54.
A.
B.
C.
D.
 H2 O(l) 
H3 O (aq)
CN (aq)

2.2x105M 2.2x107M
HCl
HCN
H3 O
NaCN
Which of the following equations correctly represents the reaction of a metallic oxide
with water?
K 2O  H2O  2KOH
SO3  H2O  H2SO4
Na2O  H2O  Na2O2  H2
NaOH  H2O  Na  2H2O
55.
A definition for a Brönsted-Lowry acid should contain which of the following phrases?
A. the donation of H
B. the donation of OH
C. the acceptance of H
D. the acceptance of OH
56.
Which equation represents the reaction of a Brönsted-Lowry base with water?
A. 2Na  2H2O  2NaOH  H2
B. N2H4  H2O  N2H5   OH
C. HPO42  H2O  H3 O  PO43
D. H2C2O4  H2O  H3 O  HC2O4 
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57.
Given the equilibrium…
H2BO3   H2PO4   H3BO3  HPO42
Which is the strongest acid?
A. HPO42
B. H3BO3
C. H2PO4 
D. H2BO3 
58.
A.
B.
C.
D.
Which species will result in a solution with the greatest H3O  ?
NaCN
Na3PO4
Na2 CO3
Na2 C2 O4
59.
Which species is not amphiprotic?
A. H2 O
B. H3BO3
C. H2PO4 
D. H2 C6H5 O7 
60.
At a given temperature a sample of pure water has a pH = 7.10. Which of the
following is true?
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61.
Which of the following is a definition of pH?
A. pH  log H3 O 
B. pH  pOH  14
C. pH  log H3 O 
D. pH  pOH  pK w
62.
What is the mass of NaOH required to prepare 100.0 mL of NaOH(aq) that has a pH
of 13.62?
A. 0.38 g
B. 0.42 g
C. 1.67 g
D. 2.40x1014 g
63.
Which of the following hypothetical acids would have the lowest conductivity?
64.
What is the net ionic equation for the hydrolysis of NH4 Cl ?
A. NH4 Cl(aq)  NH4 (aq)  Cl (aq)
B. Cl (aq)  H2O(l)  HCl(aq)  OH (aq)
C. NH4 (aq)  H2O(l)  H3 O (aq)  NH3 (aq)
D. NH4  (aq)  H2O(l)  HNH42 (aq)  OH (aq)
65.
A. 1.0
B. 5.0
C. 7.0
D. 9.0
What is the approximate pH of a 0.1 M solution of the salt NH4 Cl ?
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66.
Consider the following indicator equilibrium…
HIn(aq)
colourless
 H2 O  H3 O (aq) 
In (aq)
blue
What is the effect of adding HCl to a blue sample of this indicator?
67.
An indicator has a K a  4x106 . Which of the following is true for this indicator?
68.
Oxalic acid dihydrate is a pure, stable, crystalline substance. Which of the following
describes one of its uses in acid-base titrations?
A. buffer
B. primary standard
C. chemical indicator
D. stoichiometric indicator
69.
What is the net ionic equation that describes the reaction of HCl(aq) with
Pb  OH2 (s) ?
A. H (aq)  OH (aq)  H2O(l)
B. 2HCl(aq)  Pb  OH2 (s)  PbCl2 (s)  2H2O(l)
C. 2H (aq)  2Cl (aq)  Pb  OH2 (s)  PbCl2 (s)  2H2O(l)
D. 2H (aq)  2Cl (aq)  Pb2 (aq)  OH (aq)  Pb2 (aq)  2Cl- (aq)  2H2O(l)
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 17 / 69
70.
Which of the following would be used to prepare an acidic buffer solution?
A. HF and H3 O
B. H2 S and NaHS
C. NH3 and NH4 Cl
D. HNO3 and NaNO3
71.
Four samples of rain are collected from different geographic regions and the pH is
measured for each sample.
Which of the above samples would be classified as acid rain?
A. 1 only
B. 1 and 2
C. 1, 2 and 3
D. 1, 2, 3 and 4
72.
Which of the following is a characteristic that is common to bases?
A. They react with metals to produce OH- .
B. They produce a yellow colour in bromthymol blue solution.
C. They produce solutions with OH-  smaller than 1x10-7 .
D. They produce solutions with H3 O+  smaller than 1x10-7 .
73.
What is a common substance found in solid drain cleaner?
A. Na
B. HCl
C. NaCl
D. NaOH
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74.
Which equation contains a Brönsted-Lowry acid-base pair?
A. 2H2 (g)+O2 (g)  2H2O(g)
B. NaOH(aq)  Na+ (aq)+OH- (aq)
C. Zn(s)+2HCl(aq)  H2 (g)+ZnCl2 (aq)
D. H2 O(l)+HCN(aq)  H3 O+ (aq)+CN- (aq)
75.
Which solution would have the greatest electrical conductivity?
A. 0.1 M
B. 0.1 M
C. 0.1 M
D. 0.1 M
76.
A.
B.
C.
D.
HCl
NH3
H3BO3
CH3 OH
Which acid has the strongest conjugate base?
H2 O2
H2 CO3
HCO3
HC2 O4
77.
What is the equilibrium expression for the water ionization constant?
A. K w  Ka  Kb
B. K w =1.0x10-14
C. K w =pH+pOH
D. K w = H3 O+  OH- 
78.
A.
B.
C.
D.
What is the H3 O+  in 100.0 mL of 0.0050 M NaOH ?
5.0x10-17M
2.0x10-13M
2.0x10-12M
2.0x10-11M
79.
What is the pH of a 2.5 M KOH solution?
A. -0.40
B. 0.40
C. 13.60
D. 14.40
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80.
What is the value of K b for H2PO4- ?
A. 1.3x10-12
B. 6.2x10-8
C. 1.6x10-7
D. 7.5x10-3
81.
One of the products of the reaction between HCl(aq) and NH4 OH(aq) undergoes
hydrolysis. What is the net ionic equation for this hydrolysis reaction?
A. NH4 Cl(aq)  NH4+ (aq)+Cl- (aq)
B. Cl- (aq)+H2O(l)  HCl(aq)+OH- (aq)
C. NH4+ (aq)+H2O(l)  H3 O+ (aq)+NH3 (aq)
D. HCl(aq)+NH4 OH(aq)  NH4Cl(aq)+H2O(l)
82.
One of the species in the chemical indicator HIn exhibits a yellow colour. If acid is
added, the indicator turns red. Which of the following is correct?
83.
An indicator changes colour in the pH pH range of 6.40 to 7.20. What is the K a for
this indicator?
A. 4.0x10-7
B. 1.6x10-7
C. 0.80
D. 6.80
84.
A 25.0 mL sample of the weak acid H2 S is titrated with 31.8 mL of 0.30 M NaOH (a
strong base). What is the concentration of the acid?
A. 0.19 M
B. 0.24 M
C. 0.38 M
D. 0.76 M
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 20 / 69
85.
Which of the following is the net ionic equation for the titration reaction of NH3 (aq)
with HCl(aq) ?
A. H+ (aq)+OH- (aq)  H2O(l)
B. NH3 (aq)+H+ (aq)  NH4+ (aq)
C. NH3 (aq)+HCl(aq)  NH4Cl(aq)
D. NH3 (aq)+H+ (aq)+Cl (aq)  NH4+ (aq)+Cl (aq)
86.
What is the main function of a buffer solution?
A. to neutralize acids
B. to resist changes in pH
C. to maintain solution neutrality
D. to prevent acids from mixing with bases
87.
A.
B.
C.
D.
Which of the following would commonly be used to prepare a buffer solution?
HCl and KCl
NH3 and NH4 Cl
H2 S and Na2 SO4
Na2 CO3 and NaOH
88.
Which of the following represents a reaction that can occur between a non-metallic
oxide and water?
A. SO2 +H2O  H2SO3
B. Na2O+H2O  2NaOH
C. CaO+H2 O  Ca  OH2
D. NO2 +H2O  H2NO+O2
89.
Which of the following represents the results of tests on an acidic solution?
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 21 / 69
90.
Which of the following represents a protonated water molecule?
A. H2 O
B. H3 O
C. H2 O+
D. H3 O+
91.
Which of the following best describes a weak base?
92.
Which species will produce the greatest hydroxide ion concentration in solution?
A. NH4+
B. PO43C. HCO3D. H2 CO3
93.
Water will react most completely as an acid with...
A. SO32B. H2BO3C. C6H5 OD. CH3 COO94.
What does this equation represent?
A. K w
B. the ionization of water
C. the ion product of water
D. the equilibrium expression for water
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 22 / 69
95.
Which of the following is correct for water?
96.
What is the pH of 0.50 M LiOH ?
A. 2.0x10-14
B. -0.30
C. 0.30
D. 13.70
97.
What is the equilibrium constant expression for the predominant equilibrium in
NaHSO3 (aq) ?
H3 O+  SO3 2- 
A.
HSO3- 
OH-  H2 SO3 
B.
HSO3- 
C. NaHSO3 (aq)  Na+ (aq)+HSO3D. HSO3- (aq)+H2O(l)  H3 O+ (aq)+SO3298.
A.
B.
C.
D.
Which of the following 0.10 M solutions would have the lowest pH?
HF
NH3
HNO3
H2 CO3
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 23 / 69
99.
Which of the following describes a solution of LiClO4 (aq) with respect to hydrolysis?
A. Li+ (aq)+H2O(l)  LiOH(aq)+H+ (aq)
B. LiClO4 (aq)+H2O(l)  LiOH(aq)+HClO4 (aq)
C. ClO4- (aq)+H2 O(l)  OH- (aq)+HClO4 (aq)
D. No hydrolysis reaction occurs.
100.
A.
B.
C.
D.
Which of the following is a basic salt solution?
NH3 (aq)
NH4I(aq)
KNO3 (aq)
Na2CO3 (aq)
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 24 / 69
Written Questions
1.
Calculate the pH of a 1.5 M H2 S solution. (4 marks)
2.
Consider the following reaction…
2HCl(aq)  Ba  OH2 (s)  BaCl2 (aq)  2H2O(l)
When 3.16 g samples of Ba  OH2 were titrated to the equivalence point with an HCl solution,
the following data were recorded…
Using the data above, calculate the original HCl . (4 marks)
3.
In aqueous solutions H3 O is the strongest acid present. This phenomenon is called
the leveling effect. Explain why this occurs. (2 marks)
4.
A 1.00 M OCl solution has an OH  of 5.75x104M .
A. Calculate K b for OCl . (3 marks)
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 25 / 69
B. Calculate K a for HOCl . (1 mark)
5.
Calculate the mass of H2PO4  needed to prepare 2.0 L of a solution with a pH of
12.00. (3 marks)
6.
Consider a Brönsted-Lowry acid-base equation, where HNO2 is a reactant and
H2PO4  is a product.
A. Complete the following equation: HNO2  ?  ?  H2PO4  (1 mark)
B. Identify the weaker base in the equilibrium in part a). (1 mark)
7.
A 0.0200 M solution of methylamine, CH3NH2 , has a pH of 11.40.
Calculate the K b for methylamine. (4 marks)
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 26 / 69
8.
A titration was performed by adding 0.115 M NaOH to a 25.00 mL sample of H2 SO4 .
Calculate the H2 SO4  from the following data. (3 marks)
9.
A sample of a weak acid was found to conduct an electric current better than a
sample of a strong acid. Explain these results in terms of ion concentration. (2 marks)
10.
Calculate the OH  of 0.10 M NH3 . (4 marks)
11.
A titration was performed by adding 0.175 M H2 C2 O4 to a 25.00 mL sample of
NaOH . The following data was collected…
Calculate the NaOH . (3 marks)
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 27 / 69
B. Explain why the pH at the equivalence point is greater than 7.0. (1 mark)
A 0.100 M solution of an unknown weak acid, HX , has a pH of 1.414.
12.
What is the K a for HX ? (4 marks)
13.
Consider the salt ammonium acetate, NH4 CH3 COO .
A. Write the equation for the dissociation of NH4 CH3 COO . (1 mark)
B. Write equations for the hydrolysis reactions which occur. (2 marks)
C. Explain why a solution of NH4 CH3 COO has a pH of 7.00. Support your answer with
calculations. (2 marks)
14.
A. Define the term Brönsted-Lowry conjugate acid-base pair. (1 mark)
B. Give an example of a conjugate acid-base pair. (1 mark)
Acid:
Base:
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 28 / 69
15.
Consider the acids HCl and HF …
A. Only one of the following reactions occurs. Complete the equation for the reaction which
does occur. (1 mark)
HCl  F  ?
HF  Cl  ?
B. For the reaction that occurs, are reactants or products favoured? Explain. (1 mark)
C. Explain why the other reaction will not occur. (1 mark)
16.
Calculate the H3 O  of 0.10 M HNO2 . (3 marks)
17.
Write the formula equation and the net ionic equation for the reaction between 0.10 M
H2 SO4 and 0.10 M Sr  OH2 . (3 marks)
Formula equation:
Net ionic equation:
18.
Consider the following Brönsted-Lowry equilibrium:
H2SO3 (aq)  HPO42 (aq)  H2PO4 (aq)  HSO3 (aq)
A. Identify the two Brönsted-Lowry acids in the above equilibrium. (1 mark)
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 29 / 69
B. Define the term conjugate acid. (1 mark)
19.
A 250.0 mL sample of HCl with a pH of 2.000 is completely neutralized with 0.200 M
NaOH .
A. What volume of NaOH is required to reach the stoichiometric point? (4 marks)
B. Write the net ionic equation for the above neutralization reaction. (1 mark)
C. If the HCl were titrated with a 0.200 M NH3 (aq) instead of 0.200 M NaOH , how would the
volume of base required to reach the equivalence point compare with the volume calculated
in part a) ? Explain your answer. (1 mark)
20.
Consider the following equilibrium…
energy  2H2 O  H3 O  OH
A. Explain how pure water can have a pH of 7.30. (2 marks)
B. Calculate the value of K W for the sample of water with a pH of 7.30. (2 marks)
21.
Identify a gas which causes acid rain, and write an equation showing this gas reacting
with water. (2 marks)
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 30 / 69
22.
Consider the 0.10 M solutions of the following two acids:
A. What can you conclude about the acids that will explain these different pH values? (1
mark)
You can conclude that acid HA:
You can conclude that acid HB:
B. Compare the volume of 0.10 M NaOH needed to neutralize equal volumes of each of
these acid samples. (1 mark)
23.
Consider a 1.0 M solution of NH4F …
A. Write both hydrolysis reactions that occur when NH4F is dissolved in water. (2 marks)
B. Will the above NH4F solution be acidic, basic, or neutral? Support your answer with
calculations. (2 marks)
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 31 / 69
24.
An indicator is often used during acid-base titrations…
A. Define the term transition point for an indicator. (1 mark)
B. Calculate the K a value for methyl red. (1 mark)
C. A mixture of indicators is made by combining equal amounts of methyl orange and
bromthymol blue. Complete the following table, showing the colour of each indicator and the
mixture at pH of 5 and pH of 9. (2 marks)
25.
Will HC2 O4 act predominantly as an acid or as a base in solution? Support your
answer with calculations. (3 marks)
26.
The two reactants in an acid-base reaction are HNO2 (aq) and HCO3- (aq) .
A. Write the equation for the above reaction. (2 marks)
B. Define the term conjugate acid-base pair. (1 mark)
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 32 / 69
C. Write the formulas for a conjugate acid-base pair for the above reaction. (1 mark)
27.
At 10.0 C, K w =2.95x10-15 for pure water.
A. Calculate the pH of water at 10.0 C. (3 marks)
B. A mixture of the indicators phenolphthalein and bromcresol green is added to the water.
What is the resulting colour of the mixture? Explain. (2 marks)
28.
At a particular temperature a 1.0 M H2 S solution has a pH of 3.75 . Calculate the
value of K a at this temperature. (4 marks)
29.
What is the main function of a buffer solution? (1 mark)
30.
The ion H2PO4- is an amphiprotic anion.
A. Define the term amphiprotic. (1 mark)
B. Write the balanced equation for the reaction when H2PO4- reacts with HF . (2 marks)
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 33 / 69
31.
Write an equation for a reaction in which H2 O acts only as a Brönsted-Lowry base. (2
marks)
32.
Calculate the pH of 0.25 M Sr  OH2 . (2 marks)
33.
Calculate the pH of 0.25 M NH4 Cl . (5 marks)
34.
A 0.1 M unknown acid is titrated with 0.10 M NaOH and the following titration curve
results…
A. Choose a suitable indicator (other than phenolphthalein) and give a reason for your choice.
(1 mark)
B. Is the unknown acid weak or strong? Explain. (2 marks)
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35.
The cyanide ion, CN- , is a Brönsted-Lowry base.
A. Define Brönsted-Lowry base. (1 mark)
B. Write the equation representing the reaction of CN- with water. (2 marks)
C. Identify a conjugate pair in B. above. (1 mark)
36.
Write an equation to show the ionization of water. (2 marks)
37.
Calculate the pH of 1.50 M NH3 . (5 marks)
38.
Consider the following buffer equilibrium...
HF(aq)
high concentration
+H2 O(l) 
H3 O+
+
F-
low concentration high concentration
Using Le Chatelier’s Principle, explain what happens to the pH of the buffer solution when a
small amount of NaOH is added.
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 35 / 69
39.
Consider the following equilibria…
A. In equation I above, the reactants are favoured. Identify the stronger acid. (1 mark)
B. In equation II above, the products are favoured. Identify the stronger acid. (1 mark)
C. Consider the following reaction… HOCN+ClO-  OCN- +HClO . Does this reaction favour
reactants or products? Explain. (2 marks)
40.
At 60 C, the pH for pure water is 6.51. Determine the value of K w at this temperature.
(3 marks)
41.
Calculate the pH of 0.35 M H2 CO3 . (4 marks)
42.
A strong acid strong base titration has a pH of 7.0 at the equivalence point. A weak
acid strong base titration has a pH greater than 7.0 at the equivalence point.
A. What causes the difference in these pH values? (2 marks)
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 36 / 69
B. Select one indicator which could be used for both titrations. (1 mark)
43.
Write a chemical reaction showing an amphiprotic anion reacting as a base in water.
(2 marks)
44.
Calculate the pOH of 0.25 M Sr  OH2 . (2 marks)
45.
A 2.00 M diprotic acid has a pH of 0.50. Calculate the K a value. (5 marks)
46.
The following two experiments were conducted…
Titration A
Titration B
A strong acid was titrated with a strong base.
A weak acid was titrated with a strong base.
A. How does the pH at the equivalence point of Titration B compare with the pH at the
equivalence point in titration A? (1 mark)
B. Explain your answer to A. (2 marks)
47.
An acid-base reaction occurs between HSO3- and IO3 - .
Write the equation for the equilibrium that results. (1 mark)
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B. Identify one conjugate acid base pair in the reaction. (1 mark)
C. State whether reactants or products are favoured and explain how you arrived at your
answer. (2 marks)
48.
At 10 C, K w =2.95x10-15 .
A. Determine the pH of water at 10 C. (3 marks)
B. State whether water at this temperature is acidic, basic or neutral, and explain. (1 mark)
49.
Calculate the pH of 0.50 M H2 S . (4 marks)
50.
Outline a procedure to prepare a buffer solution. (3 marks)
51.
Calculate the pH of a 0.60 M NH4I . Start by writing the equation for the predominant
equilibrium reaction. (5 marks)
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52.
Calculate the pH of a 0.35 M solution of the salt ammonium bromide. Begin by writing
the equation for the predominant equilibrium. (5 marks)
53.
The following three solutions are mixed in a fourth container, what pH results? (3
marks)
54.
Complete the following equilibrium, then predict whether the reactants or products will
be favoured and explain why. (3 marks)
Equilibrium equation
Reactants or products
favoured?
Explanation:
HSO3- +HSO4-  _____  _____
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55.
Calculate the initial concentration of a KF salt solution that has a pH of 8.65. Begin by
writing the equation for the predominant equilibrium reaction. (5 marks)
56.
In three separate trials, 10.00 mL samples of H2 SO4 were titrated with 0.40 M NaOH
and the results are as below. Calculate the concentration of the H2 SO4 . (3 marks)
57.
A Solution of Sr  OH2 (aq) is titrated with H2 SO4 . Explain what will happen to the
electrical conductivity during the titration. Begin by writing the balanced formula
equation, including states, to support your answer. (3 marks)
Balanced formula
equation
Explanation
58.
Aniline  C6H5NH2  is a weak base with a Kb =4.3x10-10 . Calculate the concentration
of an aniline solution that has a pH of 8.80. Begin by writing the equation for the
predominant equilibrium. (5 marks)
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 40 / 69
59.
Calculate the OH  that results when 800.0 mL of 0.010 M HCl is mixed with 1.216
g Sr  OH2 . Assume no volume change on mixing. (3 marks)
60.
For the reactant pair KHC2 O4 and Na2HPO4 , write the net ionic equation for the
predominant equilibrium reaction that will be established. Predict whether the
equilibrium will favour reactants or products and explain why. (3 marks)
Equilibrium
Reactants or products
favoured?
Explanation
R. Janssen, MSEC Chemistry 12 Provincial Workbook (Unit 04), P. 41 / 69
Multiple Choice Questions Answer Key
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
D
D
C
B
C
C
C
D
C
A
D
C
B
A
D
A
C
D
B
A
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
C
B
B
D
C
D
D
A
C
B
C
D
B
D
B
B
A
B
B
D
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
58
59
60
A
D
D
B
A
C
C
B
B
A
B
C
D
A
A
B
C
D
B
D
61
62
63
64
65
66
67
68
69
70
71
72
73
74
75
76
77
78
79
80
C
C
B
C
B
A
C
B
C
B
B
D
D
D
A
A
D
C
D
A
81
82
83
84
85
86
87
88
89
90
91
92
93
94
95
96
97
98
99
100
C
D
B
A
B
B
B
A
A
D
A
B
C
B
C
D
A
C
D
D
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Written Questions Answer Key
1.
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2.
3.
4.
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5.
6.
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7.
8.
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9.
10.
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11.
12.
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13.
14.
15.
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16.
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17.
18.
19.
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20.
21.
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22.
23.
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24.
25.
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26.
27.
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28.
29.
30.
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31.
32.
33.
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34.
35.
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36.
37.
38.
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39.
40.
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41.
42.
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43.
44.
45.
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46.
47.
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48.
49.
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50.
51.
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52.
53.
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54.
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55.
56.
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57.
58.
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59.
60.