Name ____________________________________________________ Date ______________ Class _________
Chapter 19 Study Guide: Acids, Bases & Titration – Answers
1. Differentiate between Arrhenius, Brønsted-Lowry & Lewis acids and bases.
Arrhenius
Brønsted-Lowry
Lewis
ACIDS
Form hydronium ions (H3O+)
in aqueous sol’n
Proton (H+) donors
Electron pair acceptors
BASES
Form hydroxide ions (OH-) in
aqueous sol’n
Proton (H+) acceptors
Electron pair donors
2. Label the acid (A), base (B), conjugate acid (CA) and conjugate base (CB) in each of the following
reactions:
A
B
CB
CA
a. H2CO3 + CaCl2  CaCO3 + 2HCl
A
B
CB
CA
b. 2 HCl + Ca(OH)2  CaCl2 + 2 H2O
3. Write the names or formulas for the following acids and bases:
a. H2S – hydrosulfuric acid
e. Nitrous acid – HNO2
b. HC2H3O2– acetic acid
f. Calcium hydroxide – Ca(OH)2
c. NaOH – sodium hydroxide
g. hydrochloric acid – HCl
d. HNO3 – nitric acid
Use the following formulas for the remaining questions: (these will not be given to you on
the test!)
pH + pOH = 14
Kw = [H3O+][OH-] = 1.0  10-14
+
MAVA nA = MBVB nB
pH = -log[H3O ]
pOH = -log[OH-]
4. What is the pH of 1.34 x 10-4 M solution of hydrochloric acid?
pH = -log [1.34 x 10-4M] = 3.873
5. Determine the pH and pOH of a 1.85 x 10-6 M NaOH solution.
pOH = -log [1.85 x 10-6M] = 5.733
pH = 14 – 5.73 = 8.267
6. What is the [OH-] of an aqueous solution that has a pH of 5.0?
pOH = 9
[OH-] = 1 x 10-9 M
7. What is the pH and the pOH of a solution with a volume of 2000. mL containing 35 grams of
NaOH? Is this solution acidic or basic?
35gNaOH
1molNaOH

 0.44 M NaOH
2.000L
40.00gNaOH
pOH = -log[0.44M] = 0.36 pH = 14 – 0.36 = 13.64

basic
8. What is the pH of a solution that contains 0.075 grams of hydrobromic acid in 200. mL of
water?
0.075gHBr 1molHBr

 0.0046M HBr
0.200L
80.91gHBr
pH = -log [0.0046M] = 2.34

9. What is the [H3O+] of a solution that is 1.0 x 10-3 M NaOH? What is the pH?
[H3O+] = 1.0 x 10-11 M
pH = 11
10. What volume of 1.75 M HNO3 is needed to neutralize 30.0 mL of 1.50 M LiOH?
M1 = 1.75 M
M2 = 1.5 M
V1 = ?
V2 = 30.0 mL
n1 = 1
n2 = 1
M 2V2 n2
V1 
 25.7mL HNO3
M1n1
11. 95.0 mL of 0.15 M HCl are titrated with 65.0 mL of Ca(OH)2. What is the molarity of Ca(OH)2?

M1 = 0.15 M HCl
M2 = ?
V1 = 95.0 mL
V2 = 65.0 mL
n1 = 1
n2 = 2
M1V1n1
M2 
 0.11M Ca(OH)2
V2 n2
12. What volume of 0.075 M H3PO4 is required to neutralize 45.0 mL of 0.015 M Ba(OH)2?

M1 = 0.075 M
M2 = 0.015 M
V1 = ?
V2 = 45.0 mL
n1 = 3
n2 = 2
M 2V2 n2
V1 
 6.0mL H3 PO4
Each blank can be completed with a term, short
M1n1
phrase, or a number.
19.1 Acid-Base Theories

13. three
14. Arrhenius
15. hydroxide ions
16. proton (H+)
17. acceptor
18. electron pair
19. donor
20. monoprotic
21. diprotic
22. conjugate acid-base pair
23. amphoteric
Compounds can be classified as acid or bases
according to __13__ different theories. An _14__ acid
yields hydrogen ions in aqueous solution. An Arrhenius
base yields __15_ in aqueous solution. A Brønsted-Lowry
acid is a __16__ donor. A Brønsted-Lowry base is a proton
__17__. In the Lewis theory, an acid is an __18__ acceptor.
A Lewis base is an electron-pair __19__.
An acid with one ionizable hydrogen atom is
called a __20__ acid, while an acid with two ionizable
hydrogen atoms is a called a __21__ acid.
A __22__ is a pair of substances related by the gain
or loss of a hydrogen ion. A substance that can act as
both an acid and a base is called __23__.
Classify as: always true, AT; sometimes true, ST; or never true, NT.
24. NT
Hydrochloric acid is a strong acid that is diprotic.
25. NT
The ammonium ion, NH4+, is a Brønsted-Lowry base.
26. AT
A Brønsted-Lowry base is a hydrogen-ion acceptor.
27. ST
A compound can act as both an acid and a base.
28. AT
PBr3 is a Lewis base.
Matching
29. G
monoprotic acids
30. D
triprotic acids
31. A
acid properties
32. H
base properties
33. E
conjugate base
34. I
conjugate acid
35. C
hydronium ion (H3O+)
36. F
Lewis acid
37. B
Lewis base
a. tastes sour and will change the color of an acidbase indicator
b. an electron-pair donor
c. a water molecule that gains a hydrogen ion
d. acids that contain 3 ionizable hydrogens
e. particle that remained when an acid has donated a
hydrogen ion
f. an electron-pair acceptor
g. acids that contain one ionizable hydrogen
h. tastes bitter and feels slippery
i. particle formed when a base gains a hydrogen ion
38. Identify the Lewis acid and Lewis base in the following reaction. Explain.
Lewis base: e- pair donor
Lewis acid: e- pair acceptor
19.2 Hydrogen Ions and Acidity
39. ionize
40. 1 x 10-7M
41. 0-14
42. hydrogen ion
43. acidic
44. basic
45. neutral
46. 7
47. ion-product
48. hydrogen
49. hydroxide
Each blank can be completed with a term, short phrase, or
a number.
Water molecules can __39__ to form hydrogen ions (H+)
and hydroxide ions (OH-). The concentrations of these ions in
pure water at 25oC are both equal to __40__ M.
The pH scale, which has a range from ___41_, is used to
denote the __42__ concentration of a solution. On this scale, 0
is strongly __43__, 14 is strongly __44__, and 7 is __45__. Pure
water at 25oC has a pH of __46__.
The __47__ constant for water has a value of 1.0 x 10-14.
Thus, the product of the concentrations of __48__ ions and
__49__ ions in aqueous solution will always equal 1.0 x 10-14.
Classify as: always true, AT; sometimes true, ST; or never true, NT.
50. AT
In an acidic solution, [H+] is greater than [OH-].
51. ST
pH indicators can give accurate pH readings for solutions.
52. AT
If the [H+] in a solution increases, the [OH-] must decrease.
53. NT
The [OH-] is less than 10-7 M in a basic solution.
54. NT
The definition of pH is the negative logarithm of the hydroxide-ion concentration.
Matching
55. C
alkaline solutions
a. Aqueous solution in which [H+] and [OH-]are
equal
b. Product of hydrogen ion and hydroxide ion
56. F
pH
concentrations for water
57. E
self-ionization
c. Base solutions
d. Solution in which [H+] is less than [OH-]
58. A
neutral solution
59. B
ion product constant for water (Kw) e. Reaction in which two water molecules
produce ions
60. G
acidic solution
f. The negative logarithm of the hydrogen-ion
concentration
61. D
basic solution
g. Solution in which [H+] is greater than [OH-]
62. Calculate the hydroxide-ion concentration, [OH-], for an aqueous solution in which [H+] is 1 x
10-10 M. Is this solution acidic, basic, or neutral?
[OH-] = 1 x 10-4 M basic
63. Determine the hydrogen-ion concentrations for aqueous solutions that have the following pH
values.
a. 3
1 x 10-3 M
b. 6 1 x 10-6 M
c. 10 1 x 10-10 M
19.3 Strengths of Acids and Bases
Classify as: always true, AT; sometimes true, ST; or never true, NT.
64. ST
Acids are completely dissociated in aqueous solution.
65. NT
Diprotic acids lose both hydrogens at the same time.
66. AT
Acid dissociation constants for weak acids can be calculated from experimental data.
67. ST
Bases react with water to form hydroxide ions.
a. Ratio of the concentration of the dissociated (or
ionized) form of an acid to the concentration of the
undissociated acid
69. E
weak acids
b. Bases that dissociate completely into metal ions and
70. A
acid dissociation constant (Ka)
hydroxide ions in solution
c. Acids that ionize completely in aqueous solution
71. B
strong bases
d. Bases that do not dissociate completely in aqueous
72. D
weak bases
solution
73. F
base dissociation constant (Kb) e. Acids that are only partially ionized in aqueous
solution
f. Ratio of the concentration of conjugate acid times
19.4 Neutralization Reactions
concentration of conjugate base
Matching
68. C
strong acids
74. acid
75. hydroxide
76. water
77. neutralization
78. titration
79. equivalence
80. end point
Each blank can be completed with a term, short
phrase, or a number.
In the reaction of a(n) __74__ with a base, hydrogen
ions and __75__ ions react to produce ___76_. This reaction,
called ___77_, is usually carried out by __78__. The _79__ in a
titration is the point at which the acid has been neutralized.
At the __80__ point of a titration, the number of equivalents
of acid equals the number of equivalents of base.
Classify as: always true, AT; sometimes true, ST; or never true, NT.
81. AT
A solution of known concentration is called a standard (or stock) solution.
82. AT
The endpoint of a titration of a strong base with a strong acid occurs when [H+] = [OH-].
83. ST
The point of neutralization is the end point of a titration.
84. NT
The reaction of an acid and a base produces only water.
Matching
85. C
titration
86. E
neutralization reactions
87. A
equivalence point
88. B
standard solution
89. D
end point
a. When the number of moles of hydrogen ions
equals the number of moles of hydroxide ions
b. A solution of known concentration
c. A process for determining the concentration of a
solution by adding a known amount of standard
solution
d. Point of neutralization
e. Reactions between acids and bases to produce a
salt and water
90. Complete and balance the equations for the following acid-base reactions.
a. H3PO4 + Al(OH)3  3 H2O + AlPO4
b. 2 HI + Ca(OH)2  2 H2O + CaI2
19.5 Salts in Solution
91. salt
92. acidic
93. basic
94. neutral
95. hydrolyze or react with
96. strong
97. weak
98. buffer
99. capacity
Matching
Each blank can be completed with a term, short phrase, or
a number.
A __91__ forms when an acid is neutralized by a base.
Salts can be neutral, __92__, or ___93_ in solutions. Salts of
strong acid-strong base reaction produce ___94_ solutions with
water. Salts formed from the neutralization of weak acids or
weak bases __95__ water. They produce solutions that are
acidic or basic.
For example, the pH of a solution at the equivalence
point is greater than 7 for a _96__ base-__97__ acid titration.
Solutions that resist changes in pH are called __98__ solutions.
The buffer __99__ is the amount of acid or base that can be
added to a buffer without changing the pH greatly.
f. The cations or anions of a dissociated salt remove
hydrogen ions from or donate hydrogen ions to water
g. The amount of acid or base that can be added to a buffer
solution before a significant change in pH can occur
h. The salt produced by the titration of ammonia with
hydrochloric acid
i. A solution in which the pH remains relatively constant
when small amounts of acid or base are added
100.
F
salt hydrolysis
101.
I
buffer
102.
G
buffer capacity
103.
H
NH4Cl
104.
Predict whether an aqueous solution of each salt will be acidic, basic, or neutral.
a. NH4Cl
acidic
b. Na2CO3
basic
c. NH4NO3
acidic
19.1 Acid-Base Theories
105.
Identify the hydrogen ion donor(s) and the hydrogen ion acceptor(s) for ionization of
H2SO4 in water. Label the conjugate acid-base pairs.
A
B
CB
CA
H2SO4+ H2O  HSO4 + H3O+
106.
Identify all of the ions that may be formed when H3PO4 ionizes in water.
H+, H2PO4-, HPO42-, PO43-
107.
Classify the following acids as monoprotic, diprotic, or triprotic.
a. HCOOH – monoprotic b. HBr – monoprotic c. H2SO3 – diprotic d. H3ClO4 – triprotic
108.
What would you expect to happen when lithium metal is added to water? Show the
chemical reaction.
Li + H2O  LiOH + H2
109.
In the following chemical reaction, identify the Lewis acid and base.
BF3 + F-  BF4-
19.2 Hydrogen Ions and Acidity
110.
A solution has a hydrogen ion concentration of 1 x 10-6M. What is its pH? pH = 6
111.
What is the pH of a solution if the [H+] = 7.2 x 10-9M?
pH = 8.14
112.
What is the pOH of a solution if the [OH-] = 3.5 x 10-2M?
pOH = 1.46
113.
What is the pOH of a solution that has a pH of 3.4?
pOH = 10.6
114.
Classify each solution as acidic, basic, or neutral.
a. [H+] = 2.5 x 10-9M
basic
d. [H+] = 1 x 10-7M
neutral
b. pOH = 12.0
acidic
e. pH = 0.8
acidic
c. [OH-] = 9.8 x 10-11M acidic
115.
Calculate the pH of each solution and classify each as acidic, basic or neutral.
a. [H+] = 1 x 10-5M
5, acidic
c. [OH-] = 2.2 x 10-7M
b. [H+] = 4.4 x 10-11M
10.85, basic d. pOH = 1.4
7.34, basic
12.6, basic
116.
Why is there a minus sign in the definition of pH?
117.
A solution has a pOH of 12.4. What is the pH of the solution? pH = 1.6
118.
What is the pH of a solution with [OH-] = 1 x 10-3M? pH = 11
119.
What is the pH of a 25.0 L solution containing 0.450 grams of HCl, 1.70 grams of H2SO4, and
2.50 grams of HNO3?
0.450 g HCl x 1 mol HCl/36.46 g HCl = 0.01234 mol H+
1.70 g H2SO4 x 1 mol H2SO4/98.08 g H2SO4 = 0.01733 mol H2SO4 x 2 = 0.03467 mol
H+
2.50 g HNO3 x 1 mol HNO3/63.02 g HNO3 = 0.03967 mol H+
mol H+ = 0.01234 mol H+ + 0.01733 mol H+ = 0.03967 mol H+/25.0 L = 0.00347
M H+
pH = -log(0.00347) = 2.460 = pH
120.
What is pH and pOH of a solution that was made by adding 250. mL of water to 450. mL
of 2.75 x 10-3 M HCl?
M1 = 2.75 x 10-3 M HCl
M2 = ?
V1 = 450. mL
V2 = 450. mL + 250. mL = 700. mL
M 1V1
M2 
 1.77  10 3 M HCl
V2
pH = -log[H3O+] = =log(1.77 x 10-3) = 2.753
pH = 2.753
pOH = 11.247
19.3 Strengths of Acids and Bases
121.
Rank 1M of these compounds in order of increasing hydrogen ion concentration: weak
acid, strong acid, strong base, weak base.
Strong base, weak base, weak acid, strong acid
122.
Match each solution with its correct description.
a. dilute, weak acid (4)
(1) 18M H2SO4 (aq)
b. dilute, strong base (2)
(2) 0.5M NaOH (aq)
c. concentrated, strong acid (1)
(3) 15M NH3 (aq)
d. dilute, strong acid (5)
(4) 0.1M HC2H3O2(aq)
e. concentrated, weak base (3)
(5) 0.1M HCl (aq)
123.
A 0.10M solution of formic acid has an equilibrium [H+] = 6.3 x 10-5M. What is the pH of
this solution?
pH = -log[H3O+] = -log(6.3 x 10-5) = 4.20
124.
A 0.10M solution of hydrocyanic acid, HCN, has an equilibrium hydrogen ion
concentration of 6.3 x 10-6M. What is the pH of this solution?
pH = -log[H3O+] = -log(6.3 x 10-6) = 5.20
19.4 Neutralization Reactions
125.
What is the molarity of a sodium hydroxide solution if 38 mL of the solution is titrated
to the end point with 14 mL of 0.70M sulfuric acid?
Acid – H2SO4 Base – NaOH
M1 = 0.70 M
M2 = ?
MVn
M 2  1 1 1  0.52 M NaOH
V1 = 14 mL
V2 = 38 mL
V2 n2
n1 = 2
n2 = 1
126.
If 24.6 mL of a Ca(OH)2 solution are needed to neutralize 14.2 mL of a 0.0140M
HC2H3O2, what is the concentration of the calcium hydroxide solution?
Acid – HC2H3O2
Base – Ca(OH)2
M1 = 0.0140 M
M2 = ?
M 1V1n1
M

 4.04  10 3 M Ca(OH ) 2
2
V1 = 14.2 mL
V2 = 24.6 mL
V2 n2
n1 = 1
n2 = 2
127.
A 12.4 mL solution of H2SO4 is completed neutralized by 19.8 mL of 0.0100M Ca(OH)2.
What is the concentration of the H2SO4 solution?
Acid – H2SO4
Base – Ca(OH)2
M1 = ?
M2 = 0.0100 M
M Vn
M 1  2 2 2  0.0160 M H 2 SO4
V1 = 12.4 mL
V2 = 19.8 mL
V1n1
n1 = 2
n2 = 2
128.
What volume of 0.12M Ba(OH)2 is needed to neutralize 12.2 mL of 0.25M HCl?
Acid – HCl
Base – Ba(OH)2
M1 = 0.25 M
M2 = 0.12 M
MVn
V1 = 12.2 mL
V2 = ?
V2  1 1 1  13mLBa (OH ) 2
M 2 n2
n1 = 1
n2 = 2
129.
 A 55.0-mg sample of Al(OH)3 is reacted with 0.200M HCl. How many milliliters of the
acid are needed to neutralize the Al(OH)3?
Al(OH)3 + 3HCl  3H2O + AlCl3
55.0mgAl(OH ) 3 
130.
1molAl (OH ) 3
1g
3molHCl
1Lsol ' n
1000mL




 10.6mLHCl
1000mg 78.01gAl (OH ) 3 1molAl (OH ) 3 0.200molHCl
1L
Complete the following rules:
a. strong acid + strong base  neutral solution
b. strong acid + weak base  acidic solution
c. weak acid + strong base  basic solution