Chemistry 122 – Practice Exam
Matching
Match each item with the correct statement below.
a. calorimeter
d.
b. calorie
e.
c. joule
f.
enthalpy
specific heat
heat capacity
____
1. quantity of heat needed to change the temperature of 1 g of a substance by 1 C
____
2. quantity of heat needed to change the temperature of an object by 1 C
____
3. heat content of a system at constant pressure
Match each item with the correct statement below.
a. substituent
e.
b. structural isomers
f.
c. geometric isomers
g.
d. stereoisomers
asymmetric carbon
trans configuration
cis configuration
____
4. arrangement in which substituted groups are on the same side of a double bond
____
5. arrangement in which substituted groups are on opposite sides of a double bond
____
6. compounds that differ in the orientation of groups around a double bond
____
7. carbon atom to which four different atoms or groups are attached
Match each item with the correct statement below.
a. condensed structural formula
d.
b. homologous series
e.
c. unsaturated compound
saturated compound
complete structural formula
____
8. formula showing all the atoms and bonds in a molecule
____
9. structural formula in which some bonds and/or atoms are left out
____
10. organic compound that contains at least one double or triple carbon-carbon bond
Match each item with the correct statement below.
a. substitution reaction
d.
b. addition reaction
e.
c. hydration reaction
hydrogenation reaction
dehydrogenation reaction
____
11. a reaction in which an atom or group of atoms replaces another atom or group of atoms
____
12. a reaction involving the addition of hydrogen to a carbon—carbon double bond to produce an alkane
____
13. a reaction involving the addition of water to an alkene
____
14. a reaction involving the loss of hydrogen
Multiple Choice
Identify the choice that best completes the statement or answers the question.
____
15. What happens to the energy produced by burning gasoline in a car engine?
a. The energy is lost as heat in the exhaust.
b. The energy is transformed into work to move the car.
c. The energy heats the parts of the engine.
d. all of the above
____
16. A piece of metal is heated, then submerged in cool water. Which statement below describes what happens?
a. The temperature of the metal will increase.
b. The temperature of the water will increase.
c. The temperature of the water will decrease.
d. The temperature of the water will increase and the temperature of the metal will decrease.
____
17. What would likely happen if you were to touch the flask in which an endothermic reaction were occurring?
a.
b.
c.
d.
The flask would probably feel cooler than before the reaction started.
The flask would probably feel warmer than before the reaction started.
The flask would feel the same as before the reaction started.
none of the above
____
18. If heat is released by a chemical system, an equal amount of heat will be ____.
a. absorbed by the surroundings
c. released by the surroundings
b. absorbed by the universe
d. released by the universe
____
19. Which of the following is transferred due to a temperature difference?
a. chemical energy
c. electrical energy
b. mechanical energy
d. heat
____
20. In an exothermic reaction, the energy stored in the chemical bonds of the reactants is ____.
a. equal to the energy stored in the bonds of the products
b. greater than the energy stored in the bonds of the products
c. less than the energy stored in the bonds of the products
d. less than the heat released
____
21. When your body breaks down sugar completely, how much heat is released compared to burning the same amount of
sugar in a flame?
a. The body releases more heat.
b. The body releases less heat.
c. The body releases the same amount of heat.
d. The body releases no heat.
____
22. A piece of candy has 5 Calories (or 5000 calories). If it could be burned, leaving nothing but carbon dioxide and water,
how much heat would it give off?
a. 500 calories
c. 5000 joules
b. 5 kilocalories
d. Not enough information is given.
____
23. How many joules are in 148 calories? (1 cal = 4.18 J)
a. 6.61 J
c.
b. 35.4 J
d.
____
24. What is the amount of heat required to raise the temperature of 200.0 g of aluminum by 10 C? (specific heat of
aluminum = 0.21
a.
b.
____
)
420 cal
4200 cal
c.
d.
42,000 cal
420,000 cal
25. What is the specific heat of a substance if 1560 cal are required to raise the temperature of a 312-g sample by 15 C?
a.
c.
0.033
0.99
b.
____
148 J
619 J
d.
0.33
1.33
26. How many kilocalories of heat are required to raise the temperature of 225 g of aluminum from 20 C to 100 C?
(specific heat of aluminum = 0.21
a.
b.
0.59 kcal
3.8 kcal
)
c.
d.
85 kcal
none of the above
____
27. The heat capacity of an object depends in part on its ____.
a. mass
c. shape
b. enthalpy
d. potential energy
____
28. Which of the following is a valid unit for specific heat?
a.
c.
b.
____
cal
d.
C
29. When 45 g of an alloy, at 25 C, are dropped into 100.0 g of water, the alloy absorbs 956 J of heat. If the final
temperature of the alloy is 37 C, what is its specific heat?
a.
b.
____
c.
0.423
d.
1.77
30. The specific heat of silver is 0.24
9.88
48.8
. How many joules of energy are needed to warm 4.37 g of silver from 25.0 C to
27.5 C?
a. 2.62 J
b. 0.14 J
c.
d.
45.5 J
0.022 J
____
31. Which of the following substances has the highest specific heat?
a. steel
c. alcohol
b. water
d. chloroform
____
32. By what quantity must the heat capacity of an object be divided to obtain the specific heat of that material?
a. its mass
c. its temperature
b. its volume
d. its energy
____
33. The amount of heat transferred from an object depends on which of the following?
a. the specific heat of the object
c. the mass of the object
b. the initial temperature of the object
d. all of the above
____
34. On what principle does calorimetry depend?
a. Hess's law
b. law of conservation of energy
c.
d.
law of enthalpy
law of multiple proportions
____
35. How can the enthalpy change be determined for a reaction in an aqueous solution?
a. by knowing the specific heat of the reactants
b. by mixing the reactants in a calorimeter and measuring the temperature change
c. by knowing the mass of the reactants
d. The enthalpy change for this type of reaction cannot be determined.
____
36. Calculate the energy required to produce 7.00 mol Cl O
2Cl (g) + 7O (g) + 130 kcal
a. 7.00 kcal
b. 65 kcal
____
2Cl O (g)
c.
d.
130 kcal
455 kcal
37. What is the standard heat of reaction for the following reaction?
Zn(s) + Cu
(aq)
Zn
(aq) + Cu(s)
( H for Cu = +64.4 kJ/mol;
a. 216.8 kJ released per mole
b. 88.0 kJ released per mole
____
on the basis of the following balanced equation.
H
for Zn
= –152.4 kJ/mol)
c. 88.0 kJ absorbed per mole
d. 216.8 kJ absorbed per mole
38. Calculate H for the following reaction.
C H (g) + H (g)
C H (g)
( H for C H (g) = 52.5 kJ/mol;
a. –137.2 kJ
b. –32.2 kJ
H
for C H (g) = –84.7 kJ/mol)
c. 32.2 kJ
d. 137.2 kJ
____
39. What is the heat of solution?
a. the amount of heat required to change a solid into a liquid
b. the amount of heat absorbed or released when a solid dissolves
c. the amount of heat required to change a vapor into a liquid
d. the amount of heat released when a vapor changes into a liquid
____
40. The H
is ____.
a. always negative
b. always positive
c. sometimes positive, sometimes negative
d. always 0
= –445.1 kJ/mol) dissolves in 10 L of water, how much heat is released?
c. 11.1 J
d. 11.1 kJ
____
41. When 1.0 g of solid NaOH ( H
a. 445.1 kJ
b. 405.1 kJ
____
42. When 10 g of diethyl ether is converted to vapor at its boiling point, about how much heat is absorbed? (C H
O,
= 15.7 kJ/mol, boiling point: 34.6 C)
a.
b.
____
____
____
2 kJ
2J
c.
d.
43. Hess's law ____.
a. makes it possible to calculate H for complicated chemical reactions
b. states that when you reverse a chemical equation, you must change the sign of
c. determines the way a calorimeter works
d. describes the vaporization of solids
H
44. Using a table that lists standard heats of formation, you can calculate the change in enthalpy for a given chemical
reaction. The change in enthalpy is equal to ____.
a.
H of products minus H of reactants
b.
H
of products plus
c.
H
of reactants minus
d.
H
of products divided by
H
of reactants
H
of products
H
of reactants
45. Calculate H for the reaction of sulfur dioxide with oxygen.
2SO (g) + O (g)
2SO (g)
( H SO (g) = –296.8 kJ/mol;
a. –98.9 kJ
b. –197.8 kJ
____
0.2 kJ
Not enough information is given.
H SO (g) = –395.7 kJ/mol)
c. 197.8 kJ
d. Not enough information is given.
46. Which expression represents a reaction rate?
a. time/mass
b. number/time
c.
d.
energy/time
time/energy
____
47. Activation energy is ____.
a. the heat released in a reaction
b. an energy barrier between reactants and products
c. the energy given off when reactants collide
d. generally very high for a reaction that takes place rapidly
____
48. Why does a higher temperature cause a reaction to go faster?
a. There are more collisions per second only.
b. Collisions occur with greater energy only.
c. There are more collisions per second and the collisions are of greater energy.
d. There are more collisions per second or the collisions are of greater energy.
____
49. Why does a catalyst cause a reaction to proceed faster?
a. There are more collisions per second only.
b. The collisions occur with greater energy only.
c. The activation energy is lowered only.
d. There are more collisions per second and the collisions are of greater energy.
____
50. What happens to a catalyst in a reaction?
a. It is unchanged.
b. It is incorporated into the products.
c.
d.
It is incorporated into the reactants.
It evaporates away.
____
51. The rate of a chemical reaction normally ____.
a. decreases as temperature increases
b. is slowed down by a catalyst
c. increases as reactant concentration increases
d. decreases as reactant concentration increases
____
52. Which of the following substances act as catalysts in the body?
a. carbohydrates
c. lipids
H
b.
nucleic acids
d.
enzymes
____
53. At equilibrium, what is the rate of production of reactants compared with the rate of production of products?
a. much higher
c. the same
b. higher
d. lower
____
54. If a reaction is reversible, what are the relative amounts of reactant and product at the end of the reaction?
a. no reactant; all product
b. no product; all reactant
c. some product; some reactant
d. The relationship between reactants and products cannot be determined.
____
55. If sulfur dioxide and oxygen can be made into sulfur trioxide, what is the reverse reaction?
a. 2SO  2SO + O
c. 2SO + O  2SO
b.
SO
+ O  SO
d.
SO
+ 2SO  3S + 4O
____
56. What happens to a reaction at equilibrium when more reactant is added to the system?
a. The reaction makes more products.
c. The reaction is unchanged.
b. The reaction makes more reactants.
d. The answer cannot be determined.
____
57. In an endothermic reaction at equilibrium, what is the effect of raising the temperature?
a. The reaction makes more products.
c. The reaction is unchanged.
b. The reaction makes more reactants.
d. The answer cannot be determined.
____
58. What is the equilibrium constant for the following reaction?
C+O
CO
a.
c.
b.
____
d.
59. In an equilibrium reaction with a K
a.
b.
reactants are favored
reaction is spontaneous
of 1
10 , the ____.
c.
d.
the products are favored
reaction is exothermic
____
60. The amount of disorder in a system is measured by its ____.
a. activation energy
c. equilibrium position
b. entropy
d. K
____
61. If a system is left to change spontaneously, in what state will it end?
a. the same state in which it began
b. the state with lowest possible energy
c. the state with the maximum disorder
d. the state with the lowest possible energy consistent with the state of maximum disorder
____
62. Which reaction results in the greatest increase in entropy?
a. A  B
c. 2A  B
b. A  2B
d. 3A  B
____
63. Which of the following statements explains why the melting of ice is a spontaneous reaction at room temperature and
pressure?
a. Melting is accompanied by a decrease of entropy.
b. Melting is accompanied by an increase of entropy.
c. Melting is accompanied by a decrease of energy.
d. Melting is accompanied by an increase of energy.
____
64. The melting of ice at temperatures above 0 C ____.
a. liberates heat
c.
b. is not spontaneous
d.
____
is not favorable
is endothermic
65. When an acid reacts with a base, what compounds are formed?
a. a salt only
c. metal oxides only
b. water only
d. a salt and water
____
66. What is a property of a base?
a. bitter taste
b. watery feel
c.
d.
strong color
unreactive
____
67. What is an acid according to Arrhenius?
a. a substance that ionizes to yield protons in aqueous solution
b. a substance that is a hydrogen ion donor
c. a substance that accepts an electron pair
d. a substance that is a hydrogen ion acceptor
____
68. Which of these is an Arrhenius base?
a. LiOH
b.
____
NH
69. What type of acid is sulfuric acid?
a. monoprotic
b. diprotic
c.
H PO
d.
CH COOH
c.
d.
triprotic
none of the above
____
70. Which compound can act as both a Brønsted-Lowry acid and a Brønsted-Lowry base?
a. water
c. sodium hydroxide
b. ammonia
d. hydrochloric acid
____
71. Which of the following reactions illustrates amphoterism?
a. H O + H O
c. HCl + H O
H O + OH
b. NaCl
d. NaOH
Na + OH
Na
____
72. What are the acids in the following equilibrium reaction?
CN + H O
HCN + OH
a. CN , H O
c. CN , OH
b. H O, HCN
d. H O, OH
____
73. What is the charge on the hydronium ion?
a. 2–
b. 2–
c.
d.
H O
+ Cl
+ OH
0
1+
____
74. If the hydrogen ion concentration of a solution is 10
M, is the solution acidic, alkaline, or neutral?
a. acidic
c. neutral
b. alkaline
d. The answer cannot be determined.
____
75. What is pH?
a. the negative logarithm of the hydrogen ion concentration
b. the positive logarithm of the hydrogen ion concentration
c. the negative logarithm of the hydroxide ion concentration
d. the positive logarithm of the hydroxide ion concentration
____
76. What characterizes a strong acid or base?
a. polar covalent bonding
b. complete ionization in water
c. ionic bonding
d. presence of a hydroxide or hydrogen ion
____
77. What is another name for the acid dissociation constant?
a. equilibrium constant
c. rate constant
b. ionization constant
d. mole fraction
____
78. A 0.12M solution of an acid that ionizes only slightly in solution would be termed ____.
a. concentrated and weak
c. dilute and weak
b. strong and dilute
d. concentrated and strong
____
79. A substance with a K of 1
a. strong acid
b. weak acid
10
would be classified as a ____.
c. strong base
d. weak base
____
80. The ionization constant (K ) of HF is 6.7
a. [HF] is greater than [H ][F ].
b.
. Which of the following is true in a 0.1M solution of this acid?
c. [HF] is equal to [H ][F ].
d.
[HF] is less than [H ][F ].
10
10
[HF] is equal to [H ][F
].
____
81. If an acid has a K = 1.6
a. acidic
b. basic
, what is the acidity of the solution?
c. neutral
d. The answer cannot be determined.
____
82. In a titration, when the number of moles of hydrogen ions equals the number of moles of hydroxide ions, what is said to
have happened?
a. The equivalence point has been reached.
b. The end point has been reached.
c. The point of neutralization has been reached.
d. The titration has failed.
____
83. How many covalent bonds can each carbon atom form?
a. 1
c. 3
b. 2
d. 4
____
84. The name for an alkyl group that contains two carbon atoms is ____.
a. diphenyl
c. dimethyl
b. ethyl
d. ethane
____
85. What is the general formula for a straight-chain alkane?
a. C H
c. C H
b. C H
d. C H
____
86. What is the condensed structural formula for 2,2-dimethylbutane?
a. CH (CH ) CH
c. (CH ) CCH CH
b. CH CH CH CH CH
d. C H (CH )
____
87. What is the name of the compound CH CH(CH )C(CH ) ?
a. 2,2,3-trimethylbutane
c. 1,1,1,2-tetramethylpropane
b. tetramethylpropane
d. isoheptane
____
88. Why are the molecules of hydrocarbons nonpolar?
a. The intermolecular attractions are strong.
b. All the bonds are single covalent bonds.
c. The electron pair is shared almost equally in all the bonds.
d. Van der Waals forces overcome polarity.
____
89. Which of the following compounds is an unsaturated hydrocarbon?
a. methane
c. nonane
b. propyne
d. methyl
____
90. Which of these compounds is an alkene?
a. methane
b. nonene
c.
d.
butyne
propanone
____
91. Which of the following compounds is a structural isomer of butane?
a. 2-methylbutane
c. 2-methylpropane
b. 2,2-dimethylbutane
d. 2,2-diethylpropane
____
92. Which of the following is true about structural isomers?
a. Structural isomers have the same molecular formula.
b. Structural isomers have different physical and chemical properties.
c. Structural isomers have the same elemental composition.
d. all of the above
____
93. A structural isomer of hexane is ____.
a. 2,2-dimethylbutane
b. cyclohexane
c.
d.
benzene
2-methylpentene
____
____
____
94. Alkanes do not have geometric isomers because the carbon atoms in their carbon-carbon bonds are ____.
a. double bonds
c. free to rotate
b. quite polar
d. asymmetric
95. How many different atoms or groups are attached to an asymmetric carbon?
a. 2
c. 6
b. 4
d. 8
96. What is the name of the functional group in the following compound?
a.
b.
____
c.
d.
isopropyl bromide
tert-butyl bromide
c.
d.
isobutyl bromide
sec-butyl bromide
98. What is the carbon skeleton of the product formed in the following reaction?
C H + HBr 
a.
c.
b.
d.
CCCBr
____
carbonyl
carboxylic acid
97. What is the common name of the following compound?
a.
b.
____
halogen
ester
CCBrC
99. In an addition reaction, which bond of the reactant is broken?
a. carbon—carbon single bond
c. carbon—carbon double bond
b. carbon—hydrogen single bond
d. carbon—hydrogen double bond
____ 100. What type of compound is CH OCH CH CH ?
a. alcohol
c. ether
b. aldehyde
d. ketone
____ 101. Name the following compound.
CH CH CH CH OC H
a. cyclohexylbutyl ether
b. butylcyclohexyl ether
c.
d.
phenylbutyl ether
butylphenyl ether
c.
d.
4-butanone
4-pentanone
____ 102. What is the name of the following compound?
a.
b.
2-butanone
2-pentanone
____ 103. Which carbon skeleton represents a ketone?
a.
b.
c.
d.
____ 104. A ketone has the general structure ____________.
a.
ROR
b.
c.
d.
____ 105. What is the name of the following compound?
a.
b.
butane
butanal
c.
d.
butanol
butanone
____ 106. Which carbon skeleton contains a carboxyl group?
a.
c.
CCCO
b.
d.
____ 107. Which carbon skeleton represents an ester?
a.
CCCCCOCC
b.
c.
d.
____ 108. When an oxygen atom is attached to a carbon atom, the carbon atom becomes more ____.
a. oxidized
c. acidic
b. reduced
d. basic
Short Answer
109. How many joules are there in 215 calories? (1 cal = 4.184 J)
110. How much heat is required to raise the temperature of 5.5
0.21
10
g of aluminum by 10 C? (specific heat of aluminum =
)
111. If 500 g of iron absorbs 22,000 cal of heat, what will be the change in temperature? (specific heat of iron = 0.11
)
112. A 55.0-g piece of copper wire is heated, and the temperature of the wire changes from 19.0 C to 86.0 C. The amount of
heat absorbed is 343 cal. What is the specific heat of copper?
113. The specific heat capacity of graphite is 0.71
. Calculate the energy required to raise the temperature of 750 g of
graphite by 160 C.
114. It takes 770 joules of energy to raise the temperature of 50.0 g of mercury by 110 C. What is the specific heat of
mercury?
115. When 64.0 g of methanol (CH OH) is burned, 1454 kJ of energy is produced. What is the heat of combustion for
methanol?
116. How much heat is required to melt 1.6 moles of NaCl ( H
= 30.2 kJ/mol) at its melting point?
117. Suppose a substance has a heat of fusion equal to 45 cal/g and a specific heat of 0.75
in the liquid state. If 5.0 kcal
of heat are applied to a 50-g sample of the substance at a temperature of 24 C, what will its new temperature be? What
state will the sample be in? (melting point of the substance = 27 C; specific heat of the solid = 0.48
; boiling point
of the substance = 700 C)
118. Consider a 67-g chunk of ice ( H
= 6.0 kJ/mol) in a beaker immersed in a water bath. To produce just enough heat to
melt the ice, how many moles of solid NaOH ( H
= –445.1 kJ/mol) must you dissolve in the water bath?
119. If you supply 36 kJ of heat, how many moles of ice at 0 C can be melted, heated to its boiling point, and completely
boiled away? ( H
= 40.5 kJ/mol;
H
120. Use the information below to calculate
2NO (g)  N O (g)
2N (g) + 2O (g)  2NO (g)
H
H
= 6.0 kJ/mol; specific heatwater = 0.0753
)
for the following reaction.
= 67.7 kJ
N (g) + 2O (g)  N O (g)
H
= 9.7 kJ
121. What is the equilibrium constant for the following reaction?
Si + O
SiO
122. What is the ion-product constant for water?
123. If the [H ] in a solution is 1
10
mol/L, what is the [OH ]?
124. Calculate the hydrogen-ion concentration [H ] for an aqueous solution in which [OH ] is 1
solution acidic, basic, or neutral?
125. A liter of impure water has 10
water?
10
mol/L. Is this
mol of hydroxide ions. What is the concentration of hydronium ions in this sample of
126. If the hydroxide-ion concentration is 1
10
M, what is the pH of the solution?
127. If the hydrogen-ion concentration is 1
10
M, what is the pOH of the solution?
128. What is the hydrogen-ion concentration if the pH is 3.7?
129. A 0.500M solution of a weak acid, HX, is only partially ionized. The [H ] was found
to be 4.02
10
M. Find the dissociation constant for this acid.
130. What is the acid dissociation constant of a weak acid if a concentration of 0.3M gives a hydrogen-ion concentration of
0.001M?
131. Calculate the acid dissociation constant of a weak monoprotic acid if a 0.5M solution of this acid gives a hydrogen-ion
concentration of 0.000 1M?
132. Write the general structure for aldehyde compounds.
133. What is the expected product when the following compound is oxidized?
CH CH CH CH OH
Numeric Response
134. If the hydrogen ion concentration is 10
135. If the hydroxide ion concentration is 10
136. If [OH ] = 1
10
M, what is the pH of the solution?
M, what is the pH of the solution?
M, what is the pH of the solution?
137. What is the pH of a solution with a concentration of 0.01M hydrochloric acid?
138. What is the pH when the hydrogen ion concentration is 7.0
10
M?
139. What percent of the composition of natural gas is methane?
Essay
140. Explain the distinction between heat capacity and specific heat. Provide an example to illustrate this distinction.
141. Describe the parts of a calorimeter and the function of each part.
142. What is entropy? Give several examples.
143. Compare and contrast the properties of acids and bases.
144. What are acids and bases according to the Brønsted-Lowry theory?
145. What are acids and bases according to the Lewis theory? Give examples.
146. How is strength different from concentration for acids and bases? Give an example.
147. Explain why carbon is able to form such a large number of compounds.
148. Describe in your own words what the difference is between unsaturated and saturated hydrocarbons. What is a saturated
compound saturated with?
149. Explain how geometric isomers differ from each other. Describe the difference between the trans and cis configurations
of geometric isomers. Provide an example of each configuration for a molecule that has geometric isomers.
150. Why is burning coal a major source of pollution?
151. Describe what happens in a substitution reaction. Give an example of a substitution reaction and name the atoms
involved in the replacement.
d
Answer Section
MATCHING
1. ANS:
OBJ:
2. ANS:
OBJ:
3. ANS:
OBJ:
E
PTS: 1
DIF: L1
REF:
17.1.3 Identify the units used to measure heat transfer.
F
PTS: 1
DIF: L1
REF:
17.1.3 Identify the units used to measure heat transfer.
D
PTS: 1
DIF: L1
REF:
17.2.1 Describe how calorimeters are used to measure heat flow.
4. ANS:
OBJ:
5. ANS:
OBJ:
6. ANS:
OBJ:
7. ANS:
OBJ:
G
22.3.2
F
22.3.2
C
22.3.2
E
22.3.3
p. 508
p. 508
p. 511
PTS: 1
DIF: L1
REF: p. 705
Describe the conditions under which geometric isomerism is possible.
PTS: 1
DIF: L1
REF: p. 705
Describe the conditions under which geometric isomerism is possible.
PTS: 1
DIF: L1
REF: p. 705
Describe the conditions under which geometric isomerism is possible.
PTS: 1
DIF: L1
REF: p. 705
Identify optical isomers.
8. ANS:
OBJ:
9. ANS:
OBJ:
10. ANS:
OBJ:
E
PTS: 1
DIF: L1
REF: p. 696
22.1.2 Define and describe alkanes.
A
PTS: 1
DIF: L1
REF: p. 696
22.1.2 Define and describe alkanes.
C
PTS: 1
DIF: L1
REF: p. 702
22.2.1 Describe the difference between unsaturated and saturated hydrocarbons.
11. ANS:
OBJ:
12. ANS:
OBJ:
13. ANS:
OBJ:
14. ANS:
OBJ:
A
23.1.3
D
23.2.3
C
23.2.3
E
23.3.4
PTS: 1
DIF: L1
REF: p. 728
Describe how halocarbons can be prepared.
PTS: 1
DIF: L1
REF: p. 734
Name the reactions of alkenes that may be used to introduce functional groups.
PTS: 1
DIF: L1
REF: p. 734
Name the reactions of alkenes that may be used to introduce functional groups.
PTS: 1
DIF: L1
REF: p. 743
Explain how dehydration is an oxidation reaction.
MULTIPLE CHOICE
15. ANS:
OBJ:
16. ANS:
OBJ:
17. ANS:
OBJ:
18. ANS:
OBJ:
19. ANS:
OBJ:
20. ANS:
OBJ:
21. ANS:
D
17.1.1
D
17.1.1
A
17.1.1
A
17.1.1
D
17.1.1
B
17.1.1
C
PTS: 1
DIF: L1
Explain how energy, heat, and work are related.
PTS: 1
DIF: L1
Explain how energy, heat, and work are related.
PTS: 1
DIF: L1
Explain how energy, heat, and work are related.
PTS: 1
DIF: L1
Explain how energy, heat, and work are related.
PTS: 1
DIF: L1
Explain how energy, heat, and work are related.
PTS: 1
DIF: L2
Explain how energy, heat, and work are related.
PTS: 1
DIF: L1
REF:
p. 505
REF:
p. 506
REF:
p. 506
REF:
p. 506
REF:
p. 506
REF:
p. 506
REF:
p. 507
OBJ: 17.1.2 Classify processes as either exothermic or endothermic.
22. ANS: B
PTS: 1
DIF: L2
REF: p. 507
OBJ: 17.1.2 Classify processes as either exothermic or endothermic.
23. ANS: D
PTS: 1
DIF: L1
REF: p. 507
OBJ: 17.1.3 Identify the units used to measure heat transfer.
24. ANS: A
PTS: 1
DIF: L1
REF: p. 508
OBJ: 17.1.3 Identify the units used to measure heat transfer.
25. ANS: B
PTS: 1
DIF: L1
REF: p. 509 | p. 510
OBJ: 17.1.3 Identify the units used to measure heat transfer.
26. ANS: B
PTS: 1
DIF: L1
REF: p. 508
OBJ: 17.1.3 Identify the units used to measure heat transfer.
27. ANS: A
PTS: 1
DIF: L1
REF: p. 508
OBJ: 17.1.3 Identify the units used to measure heat transfer.
28. ANS: A
PTS: 1
DIF: L1
REF: p. 508
OBJ: 17.1.3 Identify the units used to measure heat transfer.
29. ANS: B
PTS: 1
DIF: L2
REF: p. 509
OBJ: 17.1.3 Identify the units used to measure heat transfer.
30. ANS: A
PTS: 1
DIF: L2
REF: p. 509 | p. 510
OBJ: 17.1.3 Identify the units used to measure heat transfer.
31. ANS: B
PTS: 1
DIF: L2
REF: p. 509 | p. 510
OBJ: 17.1.3 Identify the units used to measure heat transfer.
32. ANS: A
PTS: 1
DIF: L2
REF: p. 509 | p. 510
OBJ: 17.1.3 Identify the units used to measure heat transfer. | 17.1.4 Distinguish between heat capacity and specific
heat.
33. ANS: D
PTS: 1
DIF: L1
REF: p. 512
OBJ: 17.2.1 Describe how calorimeters are used to measure heat flow.
34. ANS: B
PTS: 1
DIF: L2
REF: p. 511
OBJ: 17.2.1 Describe how calorimeters are used to measure heat flow.
35. ANS: B
PTS: 1
DIF: L2
REF: p. 512
OBJ: 17.2.1 Describe how calorimeters are used to measure heat flow.
36. ANS: D
PTS: 1
DIF: L2
REF: p. 515
OBJ: 17.2.2 Construct thermochemical equations. | 17.2.3 Solve for enthalpy changes in chemical reactions by using
heats of reaction.
37. ANS: A
PTS: 1
DIF: L2
REF: p. 516
OBJ: 17.2.2 Construct thermochemical equations.
38. ANS: A
PTS: 1
DIF: L2
REF: p. 516
OBJ: 17.2.2 Construct thermochemical equations. | 17.2.3 Solve for enthalpy changes in chemical reactions by using
heats of reaction.
39. ANS: B
PTS: 1
DIF: L1
REF: p. 525
OBJ: 17.3.2 Solve for the enthalpy change that occurs when a substance melts, freezes, boils, condenses, or dissolves.
40. ANS: C
PTS: 1
DIF: L1
REF: p. 525
OBJ: 17.3.2 Solve for the enthalpy change that occurs when a substance melts, freezes, boils, condenses, or dissolves.
41. ANS: D
PTS: 1
DIF: L2
REF: p. 526
OBJ: 17.3.2 Solve for the enthalpy change that occurs when a substance melts, freezes, boils, condenses, or dissolves.
42. ANS: A
PTS: 1
DIF: L2
REF: p. 524
OBJ: 17.3.2 Solve for the enthalpy change that occurs when a substance melts, freezes, boils, condenses, or dissolves.
43. ANS: A
PTS: 1
DIF: L1
REF: p. 527
OBJ: 17.4.1 State Hess's law of heat summation and describe how it is used in chemistry.
44. ANS: A
PTS: 1
DIF: L1
REF: p. 530
OBJ: 17.4.2 Solve for enthalpy changes by using Hess' law or standard heats of formation. | 17.2.3 Solve for enthalpy
changes in chemical reactions by using heats of reaction.
45. ANS: B
PTS: 1
DIF: L2
REF: p. 531
OBJ: 17.4.2 Solve for enthalpy changes by using Hess' law or standard heats of formation.
46. ANS:
OBJ:
47. ANS:
OBJ:
48. ANS:
OBJ:
49. ANS:
OBJ:
50. ANS:
OBJ:
51. ANS:
OBJ:
52. ANS:
OBJ:
53. ANS:
OBJ:
54. ANS:
OBJ:
55. ANS:
OBJ:
56. ANS:
OBJ:
57. ANS:
OBJ:
58. ANS:
OBJ:
59. ANS:
OBJ:
60. ANS:
OBJ:
61. ANS:
OBJ:
62. ANS:
OBJ:
63. ANS:
OBJ:
64. ANS:
OBJ:
65. ANS:
OBJ:
66. ANS:
OBJ:
67. ANS:
OBJ:
Lewis.
68. ANS:
OBJ:
Lewis.
69. ANS:
OBJ:
Lewis.
70. ANS:
OBJ:
B
18.1.1
B
18.1.1
C
18.1.2
C
18.1.2
A
18.1.2
C
18.1.2
D
18.1.2
C
18.2.1
C
18.2.1
A
18.2.1
A
18.2.2
A
18.2.2
B
18.2.3
C
18.2.3
B
18.4.2
D
18.4.2
B
18.4.2
B
18.4.2
D
18.4.3
D
19.4.1
A
19.1.1
A
19.1.2
PTS: 1
DIF: L1
REF: p. 542
Describe how to express the rate of a chemical reaction.
PTS: 1
DIF: L1
REF: p. 543
Describe how to express the rate of a chemical reaction.
PTS: 1
DIF: L1
REF: p. 545
Identify four factors that influence the rate of a chemical reaction.
PTS: 1
DIF: L1
REF: p. 546 | p. 547
Identify four factors that influence the rate of a chemical reaction.
PTS: 1
DIF: L1
REF: p. 546
Identify four factors that influence the rate of a chemical reaction.
PTS: 1
DIF: L1
REF: p. 545
Identify four factors that influence the rate of a chemical reaction.
PTS: 1
DIF: L2
REF: p. 547
Identify four factors that influence the rate of a chemical reaction.
PTS: 1
DIF: L1
REF: p. 550
Describe how the amounts of reactants and products change in a chemical system at equilibrium.
PTS: 1
DIF: L1
REF: p. 549 | p. 550
Describe how the amounts of reactants and products change in a chemical system at equilibrium.
PTS: 1
DIF: L2
REF: p. 549
Describe how the amounts of reactants and products change in a chemical system at equilibrium.
PTS: 1
DIF: L2
REF: p. 552 | p. 553
Identify three stresses that can change the equilibrium position of a chemical system.
PTS: 1
DIF: L2
REF: p. 554
Identify three stresses that can change the equilibrium position of a chemical system.
PTS: 1
DIF: L1
REF: p. 556
Explain what the value of Keq indicates about the position of equilibrium.
PTS: 1
DIF: L1
REF: p. 556
Explain what the value of Keq indicates about the position of equilibrium.
PTS: 1
DIF: L1
REF: p. 569
Describe the role of entropy in chemical reactions.
PTS: 1
DIF: L2
REF: p. 569
Describe the role of entropy in chemical reactions.
PTS: 1
DIF: L2
REF: p. 570
Describe the role of entropy in chemical reactions.
PTS: 1
DIF: L2
REF: p. 569
Describe the role of entropy in chemical reactions.
PTS: 1
DIF: L2
REF: p. 571
Identify two factors that determine the spontaneity of a reaction.
PTS: 1
DIF: L1
REF: p. 587
Define the products of an acid-base reaction.
PTS: 1
DIF: L1
REF: p. 588
Define the properties of acids and bases.
PTS: 1
DIF: L1
REF: p. 588
Compare and contrast acids and bases as defined by the theories of Arrhenius, Brønsted-Lowry, and
A
PTS: 1
DIF: L1
REF: p. 589
19.1.2 Compare and contrast acids and bases as defined by the theories of Arrhenius, Brønsted-Lowry, and
B
PTS: 1
DIF: L2
REF: p. 588
19.1.2 Compare and contrast acids and bases as defined by the theories of Arrhenius, Brønsted-Lowry, and
A
PTS: 1
DIF: L2
REF: p. 591
19.1.2 Compare and contrast acids and bases as defined by the theories of Arrhenius, Brønsted-Lowry, and
Lewis.
71. ANS:
OBJ:
Lewis.
72. ANS:
OBJ:
Lewis.
73. ANS:
OBJ:
74. ANS:
OBJ:
75. ANS:
OBJ:
76. ANS:
OBJ:
77. ANS:
OBJ:
78. ANS:
OBJ:
79. ANS:
OBJ:
80. ANS:
OBJ:
81. ANS:
OBJ:
82. ANS:
OBJ:
83. ANS:
OBJ:
84. ANS:
OBJ:
85. ANS:
OBJ:
86. ANS:
OBJ:
87. ANS:
OBJ:
88. ANS:
OBJ:
89. ANS:
OBJ:
90. ANS:
OBJ:
91. ANS:
OBJ:
92. ANS:
OBJ:
93. ANS:
OBJ:
94. ANS:
OBJ:
95. ANS:
OBJ:
A
PTS: 1
DIF: L2
REF: p. 592
19.1.2 Compare and contrast acids and bases as defined by the theories of Arrhenius, Brønsted-Lowry, and
B
PTS: 1
DIF: L2
REF: p. 591
19.1.2 Compare and contrast acids and bases as defined by the theories of Arrhenius, Brønsted-Lowry, and
D
19.2.1
B
19.2.1
A
19.2.2
B
19.3.1
B
19.3.1
C
19.3.2
B
19.3.3
A
19.3.3
D
19.3.3
A
19.4.4
D
22.1.1
B
22.1.2
D
22.1.2
C
22.1.2
A
22.1.2
C
22.1.3
B
22.2.1
B
22.2.2
C
22.3.1
D
22.3.1
A
22.3.1
C
22.3.2
B
22.3.3
PTS: 1
DIF: L1
REF: p. 594
Describe how [H+] and [OH+] are related in an aqueous solution.
PTS: 1
DIF: L1
REF: p. 595
Describe how [H+] and [OH+] are related in an aqueous solution.
PTS: 1
DIF: L1
REF: p. 596
Classify a solution as neutral, acid, or basic given the hydrogen-ion or hydroxide-ion concentration.
PTS: 1
DIF: L1
REF: p. 605
Define strong acids and weak acids.
PTS: 1
DIF: L2
REF: p. 607
Define strong acids and weak acids.
PTS: 1
DIF: L3
REF: p. 609
Describe how an acid's strength is related to the value of its acid dissociation constant.
PTS: 1
DIF: L2
REF: p. 607
Calculate an acid dissociation constant (Ka) from concentration and pH measurements.
PTS: 1
DIF: L2
REF: p. 607
Calculate an acid dissociation constant (Ka) from concentration and pH measurements.
PTS: 1
DIF: L2
REF: p. 610
Calculate an acid dissociation constant (Ka) from concentration and pH measurements.
PTS: 1
DIF: L1
REF: p. 613 | p. 614 | p. 615
Describe the relationship between equivalence point and the end point of a titration.
PTS: 1
DIF: L1
REF: p. 694
Describe the relationship between the valence electrons and bonding in carbon.
PTS: 1
DIF: L1
REF: p. 698
Define and describe alkanes.
PTS: 1
DIF: L2
REF: p. 695
Define and describe alkanes.
PTS: 1
DIF: L2
REF: p. 696 | p. 698
Define and describe alkanes.
PTS: 1
DIF: L2
REF: p. 698
Define and describe alkanes.
PTS: 1
DIF: L2
REF: p. 700
Relate the polarity of hydrocarbons to their solubility.
PTS: 1
DIF: L1
REF: p. 703
Describe the difference between unsaturated and saturated hydrocarbons.
PTS: 1
DIF: L1
REF: p. 702
Distinguish between the structures of alkenes and alkynes.
PTS: 1
DIF: L1
REF: p. 704
Explain why structural isomers have different properties.
PTS: 1
DIF: L1
REF: p. 704
Explain why structural isomers have different properties.
PTS: 1
DIF: L1
REF: p. 704
Explain why structural isomers have different properties.
PTS: 1
DIF: L2
REF: p. 705
Describe the conditions under which geometric isomerism is possible.
PTS: 1
DIF: L2
REF: p. 705
Identify optical isomers.
96. ANS:
OBJ:
97. ANS:
OBJ:
98. ANS:
OBJ:
99. ANS:
OBJ:
100. ANS:
OBJ:
101. ANS:
OBJ:
102. ANS:
OBJ:
103. ANS:
OBJ:
104. ANS:
OBJ:
105. ANS:
OBJ:
106. ANS:
OBJ:
107. ANS:
OBJ:
108. ANS:
OBJ:
B
23.1.1
B
23.1.2
B
23.1.3
C
23.2.3
C
23.2.4
D
23.2.4
B
23.3.1
B
23.3.1
B
23.3.1
D
23.3.1
C
23.3.2
C
23.3.3
A
23.3.4
PTS: 1
DIF: L1
REF: p. 726
Explain how organic compounds are classified.
PTS: 1
DIF: L1
REF: p. 727
Identify halocarbons and the IUPAC rules for naming halocarbons.
PTS: 1
DIF: L2
REF: p. 727
Describe how halocarbons can be prepared.
PTS: 1
DIF: L1
REF: p. 733
Name the reactions of alkenes that may be used to introduce functional groups.
PTS: 1
DIF: L1
REF: p. 735
Construct the general structure of an ether and describe how ethers are named.
PTS: 1
DIF: L2
REF: p. 735
Construct the general structure of an ether and describe how ethers are named.
PTS: 1
DIF: L1
REF: p. 738
Identify the structure of a carbonyl group as found in aldehydes and ketones.
PTS: 1
DIF: L1
REF: p. 737
Identify the structure of a carbonyl group as found in aldehydes and ketones.
PTS: 1
DIF: L1
REF: p. 737
Identify the structure of a carbonyl group as found in aldehydes and ketones.
PTS: 1
DIF: L2
REF: p. 738
Identify the structure of a carbonyl group as found in aldehydes and ketones.
PTS: 1
DIF: L1
REF: p. 740
Construct the general formula for carboxylic acids and explain how they are named.
PTS: 1
DIF: L2
REF: p. 741
Describe an ester.
PTS: 1
DIF: L1
REF: p. 743 | p. 744
Explain how dehydration is an oxidation reaction.
SHORT ANSWER
109. ANS:
215 cal
4.184
= 9.00
10
J
PTS: 1
DIF: L2
REF: p. 507
OBJ: 17.1.2 Classify processes as either exothermic or endothermic.
110. ANS:
Heat energy = mass specific heat temperature change
= 550 g
0.21
= 1.2
10 C
cal
PTS: 1
DIF: L2
REF: p. 508
OBJ: 17.1.3 Identify the units used to measure heat transfer.
111. ANS:
T = Temperature change =
=
PTS:
= 400 C
1
DIF:
L2
REF:
p. 508
OBJ: 17.1.3 Identify the units used to measure heat transfer.
112. ANS:
T = 86.0 C – 19.0 C = 67.0 C
specific heat =
=
= 9.31
10
PTS: 1
DIF: L2
REF: p. 509 | p. 510
OBJ: 17.1.3 Identify the units used to measure heat transfer.
113. ANS:
H = 750 g
0.71
160 C = 85,000 J
PTS: 1
DIF: L2
REF: p. 512
OBJ: 17.2.1 Describe how calorimeters are used to measure heat flow.
114. ANS:
Specific heat =
= 0.14
PTS: 1
DIF: L2
REF: p. 512
OBJ: 17.2.1 Describe how calorimeters are used to measure heat flow.
115. ANS:
H=
= 727 kJ/mol
PTS: 1
DIF: L2
REF:
OBJ: 17.2.2 Construct thermochemical equations.
116. ANS:
1.6 mol 30.2 kJ/mol = 48 kJ
p. 517
PTS: 1
DIF: L2
REF: p. 521
OBJ: 17.3.1 Classify the enthalpy change that occurs when a substance melts, freezes, boils, condenses, or dissolves.
117. ANS:
50 g
0.48
3.0 C
= 72 cal to raise the temperature of the solid to 27 C
50 g 45 cal/g = 2250 cal to melt the sample
2250 cal + 72 cal = 2322 cal
5000 cal – 2322 cal = 2678 cal remaining
T=
= 71 C
71 C + 27 C = 98 C
The substance is in a liquid state.
PTS:
OBJ:
1
DIF: L3
REF: p. 521
17.3.1 Classify the enthalpy change that occurs when a substance melts, freezes, boils, condenses, or dissolves.
118. ANS:
Heat to melt ice comes from heat released by the dissolving of NaOH.
Amount of NaOH = 67 g H O
= 0.050 mol NaOH
PTS: 1
DIF: L3
REF: p. 520 | p. 521
OBJ: 17.3.2 Solve for the enthalpy change that occurs when a substance melts, freezes, boils, condenses, or dissolves.
119. ANS:
Total heat = heat to melt ice + heat to warm water to 100 C + heat to evaporate water
Total heat = (moles ice
H ) + (moles water C T) + (moles water
H )
36 kJ = (moles of H O
6.0 kJ/mol) + (moles of H O
36 kJ = moles H O (6.0 kJ/mol + 0.0753
0.0753
100 C) + (moles of H O
100 C + 40.5 kJ/mol)
36 kJ = moles H O (54.0 kJ/mol)
moles H O =
moles H O = 0.67 mol
PTS: 1
DIF: L3
REF: p. 527
OBJ: 17.4.1 State Hess's law of heat summation and describe how it is used in chemistry.
120. ANS:
2NO (g)  N (g) + 2O (g)
H = –67.7 kJ
N (g) + 2O (g)  N O (g)
H
= 9.7 kJ
2NO (g)  N O (g)
H
= –58 kJ
PTS: 1
DIF: L3
REF: p. 528
OBJ: 17.4.2 Solve for enthalpy changes by using Hess' law or standard heats of formation.
121. ANS:
PTS: 1
DIF: L2
REF: p. 556
OBJ: 18.2.3 Explain what the value of Keq indicates about the position of equilibrium.
122. ANS:
10
M
PTS: 1
DIF: L1
REF: p. 595
OBJ: 19.2.1 Describe how [H+] and [OH+] are related in an aqueous solution.
123. ANS:
1
10
mol/L
PTS: 1
DIF: L2
REF: p. 595 | p. 596
OBJ: 19.2.1 Describe how [H+] and [OH+] are related in an aqueous solution.
124. ANS:
K
= [H ]
[OH ]
40.5 kJ/mol)
[H ] =
=
= 1 10 mol/L
The solution is acidic.
PTS: 1
DIF: L3
REF: p. 595
OBJ: 19.2.1 Describe how [H+] and [OH+] are related in an aqueous solution.
125. ANS:
10
M
PTS: 1
DIF: L1
REF: p. 598
OBJ: 19.2.2 Classify a solution as neutral, acid, or basic given the hydrogen-ion or hydroxide-ion concentration.
126. ANS:
K
= [H ]
[OH ]
[H ] =
=
=1
10- 2 mol/L
pH = - log [H ] = - log [ 1
10- 2 mol/L] = 2.0
PTS: 1
DIF: L2
REF: p. 598
OBJ: 19.2.2 Classify a solution as neutral, acid, or basic given the hydrogen-ion or hydroxide-ion concentration.
127. ANS:
K
= [H ]
[OH ]
[OH ] =
=
= 1 10- 1 mol/L
pOH = log [OH ] =
= 1.0
log [ 1
10- 1 mol/L]
PTS: 1
DIF: L2
REF: p. 601
OBJ: 19.2.2 Classify a solution as neutral, acid, or basic given the hydrogen-ion or hydroxide-ion concentration.
128. ANS:
–log [H ] = pH = 3.7
log [H ] = –3.7
[H ] = antilog(–3.7)
[H ] = 0.000 20M
PTS: 1
DIF: L2
REF: p. 600
OBJ: 19.2.2 Classify a solution as neutral, acid, or basic given the hydrogen-ion or hydroxide-ion concentration.
129. ANS:
[HX] = 0.500M – 4.02
10
M = 0.496M
K =
=
= 3.26
10
M
PTS: 1
DIF: L3
REF: p. 607
OBJ: 19.3.2 Describe how an acid's strength is related to the value of its acid dissociation constant.
130. ANS:
3
10
PTS: 1
DIF: L2
REF: p. 607
OBJ: 19.3.2 Describe how an acid's strength is related to the value of its acid dissociation constant.
131. ANS:
K =
=
=
= 0.000 000 02 = 2
10
PTS: 1
DIF: L3
REF: p. 607
OBJ: 19.3.2 Describe how an acid's strength is related to the value of its acid dissociation constant.
132. ANS:
PTS: 1
DIF: L2
REF: p. 737
OBJ: 23.3.1 Identify the structure of a carbonyl group as found in aldehydes and ketones.
133. ANS:
CH CH CH COOH
PTS:
OBJ:
1
DIF: L3
REF: p. 744 | p. 745
23.3.4 Explain how dehydration is an oxidation reaction.
NUMERIC RESPONSE
134. ANS: 7
PTS: 1
DIF: L1
REF: p. 596 | p. 597
OBJ: 19.2.2 Classify a solution as neutral, acid, or basic given the hydrogen-ion or hydroxide-ion concentration.
135. ANS: 4
PTS: 1
DIF: L1
REF: p. 601
OBJ: 19.2.2 Classify a solution as neutral, acid, or basic given the hydrogen-ion or hydroxide-ion concentration.
136. ANS: 10.0
PTS: 1
DIF: L1
REF: p. 601
OBJ: 19.2.2 Classify a solution as neutral, acid, or basic given the hydrogen-ion or hydroxide-ion concentration.
137. ANS: 2.0
PTS: 1
DIF: L2
REF: p. 596 | p. 597
OBJ: 19.2.2 Classify a solution as neutral, acid, or basic given the hydrogen-ion or hydroxide-ion concentration.
138. ANS: 2.2
PTS: 1
DIF: L2
REF: p. 597 | p. 598
OBJ: 19.2.2 Classify a solution as neutral, acid, or basic given the hydrogen-ion or hydroxide-ion concentration.
139. ANS: 80%
PTS:
OBJ:
1
DIF: L2
REF: p. 712
22.5.2 Describe the composition of natural gas and coal.
ESSAY
140. ANS:
Heat capacity is the quantity of heat required to change an object's temperature by 1 C. The heat capacity of any
particular object varies with the mass of that object (as well as with the type of material in the object). The heat capacity
of a steel girder is much greater than the heat capacity of a steel nail, for instance. Specific heat, on the other hand, does
not vary with the mass of the object, but rather, depends only on the nature of the material in the object. Specific heat is
the quantity of heat required to raise the temperature of 1 gram of a substance by 1 C. The specific heats of the steel in
the steel girder and the steel in the steel nail are identical (assuming the two steels are of the same composition). Specific
heat is a property of a particular material; heat capacity is a property of a particular object.
PTS: 1
DIF: L3
REF: p. 508
OBJ: 17.1.3 Identify the units used to measure heat transfer. | 17.1.4 Distinguish between heat capacity and specific
heat.
141. ANS:
Generally a calorimeter consists of an insulated container, water, and a temperature-measuring instrument. The insulated
container prevents heat from entering or leaving the system from the outside. There is water in the container to absorb
heat. The temperature-measuring device is often a thermometer. Some calorimeters have a stirrer to distribute the heat
evenly through the water. A bomb calorimeter may contain a set of ignition wires.
PTS: 1
DIF: L3
REF: p. 511 | p. 512
OBJ: 17.2.1 Describe how calorimeters are used to measure heat flow.
142. ANS:
Entropy is a measure of the degree of disorder in a system. A gas has more entropy than a liquid. A chemical reaction in
which there are more molecules of product than molecules of reactant will cause an increase in entropy. A solution of
sodium chloride in water has more entropy than a sodium chloride crystal.
PTS: 1
DIF: L3
REF: p. 569 | p. 570
OBJ: 18.4.2 Describe the role of entropy in chemical reactions.
143. ANS:
Both acids and bases are electrolytes; both cause indicators to change colors; and both react with each other to form
water and a salt. Acids taste sour, while bases taste bitter. Bases feel slippery. Acids react with some metals to produce
hydrogen gas.
PTS: 1
DIF: L3
REF: p. 587 | p. 588
OBJ: 19.1.1 Define the properties of acids and bases. | 19.4.1 Define the products of an acid-base reaction.
144. ANS:
According to the Brønsted-Lowry theory, acids donate protons to other substances, and bases accept protons from other
substances. Ammonia accepts a proton from water and therefore acts as a Brønsted-Lowry base. The water donates the
proton to ammonia and therefore acts as a Brønsted-Lowry acid.
PTS: 1
DIF: L3
REF: p. 590
OBJ: 19.1.2 Compare and contrast acids and bases as defined by the theories of Arrhenius, Brønsted-Lowry, and
Lewis.
145. ANS:
According to the Lewis theory, acids accept a pair of electrons to form a covalent, bond and bases donate an electron pair
to form a covalent bond. The hydrogen ion accepts electrons from the hydroxide ion to make water. The hydrogen ion is
the Lewis acid and the hydroxide ion is the Lewis base.
PTS: 1
DIF: L3
REF: p. 592
OBJ: 19.1.2 Compare and contrast acids and bases as defined by the theories of Arrhenius, Brønsted-Lowry, and
Lewis.
146. ANS:
A strong acid ionizes completely and a concentrated acid dissolves well. The same is true for strong and concentrated
bases. Ammonia is a weak base because it ionizes incompletely, but it can be concentrated because it dissolves well.
PTS: 1
DIF: L3
REF: p. 609
OBJ: 19.3.1 Define strong acids and weak acids.
147. ANS:
Because carbon has four valence electrons, each carbon atom can form four covalent bonds. Carbon-carbon bonds are
quite stable, and this fact, coupled with carbon’s capacity to form four covalent bonds, allows carbon to form long
straight or branched chains. Carbon atoms can also be joined in ring structures. Carbon can also form double and triple
carbon-carbon bonds, which further increases the number of possible molecules.
PTS: 1
DIF: L3
REF: p. 693 | p. 694
OBJ: 22.1.1 Describe the relationship between the valence electrons and bonding in carbon.
148. ANS:
Unsaturated hydrocarbons have at least one double or triple carbon-carbon bond, and saturated hydrocarbons contain
only single bonds. “Saturated” refers to the ratio of hydrogen atoms to carbon atoms. Molecules of a saturated
hydrocarbon contain the maximum number of hydrogen atoms per carbon atom.
PTS: 1
DIF: L3
REF: p. 702
OBJ: 22.2.1 Describe the difference between unsaturated and saturated hydrocarbons.
149. ANS:
Geometric isomers have atoms joined in the same order, but differ in the orientation of groups around a double bond. For
geometric isomers to exist, each carbon of the double bond must have at least one substituent. In the trans configuration,
the substituted groups are on opposite sides of the double bond. In the cis configuration, the substituted groups are on the
same side of the double bond. Examples are trans-2-butene and cis-2-butene.
PTS: 1
DIF: L3
REF: p. 705
OBJ: 22.3.2 Describe the conditions under which geometric isomerism is possible.
150. ANS:
Coal consists largely of condensed ring compounds of very high molecular mass. Due to the high proportion of these
aromatic compounds in coal, the burning of this fuel produces more soot than does the burning of the more aliphatic
fuels obtained from petroleum. In addition, the majority of the coal burned in North America contains about 7% sulfur,
which produces the major pollutants SO and SO when it burns.
PTS: 1
DIF: L3
REF: p. 715
OBJ: 22.5.2 Describe the composition of natural gas and coal.
151. ANS:
A substitution reaction is an organic reaction in which an atom or group of atoms replaces another atom or group of
atoms. An example of a replacement reaction is CH + Br  CH Br + HBr. The halogen replaces the hydrogen
because it is more reactive.
PTS:
OBJ:
1
DIF: L3
REF: p. 728
23.1.3 Describe how halocarbons can be prepared.