Download Chem 106 Final Exam Study Questions

Chem 106 Final Exam Study Questions – Spring 2017
Note: The answer key on the back contains chapter/sections from the text.
1. A student finds that the weight of an empty beaker is 12.024 g. She places a solid in the
beaker to give a combined mass of 12.108 g. To how many significant figures is the
mass of the solid known?
A) 3
B) 4
C) 2
D) 5
E) 1
2. 8.8 milliseconds is equal to how many seconds?
A) 8.8  10–2 s
B) 8.8  103 s
C) 0.88 s
D) 8.8  10–3 s
E) 8.8  102 s
3. Convert 257.4 m to decimeters.
A) 2.574  103 dm
B) 2.574  104 dm
C) 25.74 dm
D) 2.574 dm
E) none of these
4. 337.4oF is equivalent to
A) 169.7oC
B) 187.4oC
C) 205.2oC
D) 549.7oC
E) 664.9oC
5. Convert: –17.4oF = _______________ oC.
A) –88.9oC
B) 255.6oC
C) 8.1oC
D) –27.4oC
E) 26.3oC
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6. The number 2.00152 rounded to four significant figures is
A) 2.000
B) 2.002
C) 2.152
D) 2.001
E) none of these
7. A graduated cylinder contains 20.0 mL of water. An irregularly shaped object is placed
in the cylinder, and the water level rises to the 31.2-mL mark. If the object has a mass of
93.7 g, what is its density?
A) 3.00 g/mL
B) 4.68 g/mL
C) 0.120 g/mL
D) 8.37 g/mL
E) none of these
8. An object is 103.0 inches in height. Express this height in centimeters.
A) 8.58 cm
B) 105.5 cm
C) 261.6 cm
D) 0.02466 cm
E) 40.55 cm
9. Which of the following is a chemical change?
A) A damp towel dries.
B) Peanuts are crushed.
C) A “tin” can rusts.
D) Water condenses on a mirror.
E) At least two of the above (a-d) exhibit a chemical change.
10. Which of the following is an element?
A) helium
B) sugar
C) air
D) water
E) salt
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11. Which is an example of a homogeneous mixture?
A) oily water
B) soil (dust)
C) aluminum
D) vodka
E) sodium chloride
12. How many protons, electrons, and neutrons does the isotope
57
26
Fe3 have?
13. Which pair have approximately the same mass?
A) a proton and a neutron
B) a hydrogen, 11 H , and a deuterium, 21 H , atom
C) an electron and a proton
D) a neutron and an electron
14. Which of the following elements is an alkaline earth metal?
A) Na
B) Fe
C) Ca
D) Sc
E) Cu
15. The total number of atoms indicated by the formula Ca3(PO4)3 is
A) 10
B) 16
C) 7
D) 6
E) 18
16. Which of the following is an incorrect name for an acid?
A) phosphoric acid
B) sulfurous acid
C) hydrocyanic acid
D) acetic acid
E) hydrocarbonate acid
17. Give the formula for calcium hydrogen carbonate.
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18. Write the correct formula for dinitrogen pentoxide.
19. The name for NO3- is ______________.
20. Give the formula for carbon monoxide.
21. Give the formula for hypochlorous acid.
22. The name for (NH4)2SO4 is ______________.
23. The name for C2 H 3O 2 is _______________.
24. Give the formula for hydrosulfuric acid.
25. The name for (NH4)2CO3 is _______________.
26. Give the formula for chromium(III) iodide.
27. The name for Zn(OH)2 is ______________.
28. The binary compound PCl3 is called
A) phosphorus trichloride
B) monophosphorus trichloride
C) phosphorus chloride
D) triphosphorus chloride
E) none of these
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29. When the following equation is balanced using the smallest possible integers, what is
the number in front of the substance in bold type?
P4O10 + H2O  H3PO4
A) 6
B) 2
C) 10
D) 1
E) 4
30. Balance the equation
H2O2(l)  H2O(l) + O2(g)
Use the following to answer questions 31-34:
Use the following choices to classify each reaction given below (more than one choice may
apply).
a. oxidation-reduction
b. acid-base
c. precipitation
31. HNO3(aq) + NaOH(aq)  H2O(l) + NaNO3(aq)
32. HC2H3O2(aq) + CsOH(aq)  H2O(l) + CsC2H3O2(aq)
33. 2HCl(aq) + Pb(OH)2(aq)  PbCl2(s) + 2H2O(l)
34. Zr(s) + O2(g)  ZrO2(s)
Use the following to answer questions 35-37:
Use the following choices to classify each reaction given below (more than one choice may
apply).
a. oxidation-reduction
b. combustion
c. synthesis
d. decomposition
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35. 2Na(s) + H2(g)  2NaH(s)
36. C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)
37. 2GaN(s)  2Ga(s) + N2(g)
38. True or false? The formula of a compound that expresses the smallest whole-number
ratio of the atoms present is called the empirical formula.
A) True
B) False
39. Which represents the greatest number of atoms?
A) 50.0 g Zn
B) 50.0 g Fe
C) 50.0 g Cu
D) 50.0 g Al
E) all the same
40. How many atoms of calcium are present in 88.2 g of calcium?
A) 6.02  1023
B) 3.65  10–24
C) 1.33  1024
D) 5.31  1025
E) none of these
41. Calculate the mass of 13.8 moles of He.
A) 6.02  1023
B) 55.2
C) 17.8
D) 8.31  1024
E) 3.45
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42. How many moles of Ca atoms are in 580.6 g Ca?
A) 28.97 mol
B) 580.6 mol
C) 14.49 mol
D) 2.327  104 mol
E) 6.903  10–2 mol
43. Calculate the mass of 5.09  1019 molecules of HCl.
A) 8.45  10–5 g
B) 3.07  1043 g
C) 2.32  10–6 g
D) 3.08  10–3 g
E) none of these
44. A 10.6-mol sample of Co represents how many atoms?
A) 5.68  1022 atoms
B) 6.72  102 atoms
C) 1.76  10–23 atoms
D) 6.38  1024 atoms
E) 3.76  1026 atoms
45. 10.8 g of Cu represents how many moles?
A) 5.88 mol
B) 6.50  1024 mol
C) 6.86  102 mol
D) 0.170 mol
E) none of these
46. What is the molar mass of nitroglycerin, C3H5(NO3)3?
A) 185 g/mol
B) 179 g/mol
C) 227 g/mol
D) 199 g/mol
E) none of these
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47. Calculate the mass of 0.847 mol of H2SO4.
A) 5.00  1025 g
B) 1.16  102 g
C) 5.10  1023 g
D) 83.1 g
E) 8.64  10–3 g
48. A gaseous compound containing carbon and hydrogen was analyzed and found to
consist of 83.65% carbon by mass. What is the empirical formula of the compound?
A) CH
B) C7H16
C) C3H7
D) CH3
E) CH2
49. Determine the percentage composition (by mass) of oxygen in H2SO4.
A) 31.95 %
B) 78.25 %
C) 23.48 %
D) 16.31 %
E) 65.25 %
50. Choose the pair of compounds with the same empirical formula.
A) NaHCO3 and Na2CO3
B) K2CrO4 and K2Cr2O7
C) H2O and H2O2
D) C2H2 and C6H6
51. A compound contains 40.0% carbon, 6.7% hydrogen, and 53.3% oxygen (by mass).
Calculate the empirical formula.
A) C2H2O
B) C3H6O3
C) C2HO2
D) CH4O
E) CH2O
52. A compound has 40.68% carbon, 5.12% hydrogen, and 54.20% oxygen (by mass).
Calculate its empirical formula.
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53. A 2.4-mol sample of KClO3 was decomposed according to the equation
2KClO3 (s)  2KCl(s)  3O 2 (g)
How many moles of O2 are formed assuming 100% yield?
A) 2.0 mol
B) 3.6 mol
C) 1.6 mol
D) 2.4 mol
E) 1.2 mol
54. Refer to the following unbalanced equation:
C6H14 + O2  CO2 + H2O
1. Balance the equation
2. What mass of carbon dioxide (CO2) can be produced from 21.1 g of C6H14 and excess
oxygen?
A) 10.8 g
B) 32.3 g
C) 0.245 g
D) 64.7 g
E) 5.57  103 g
55. Consider the reaction
4Fe(s)  3O 2 (g)  2Fe2 O3 (s)
If 11.5 g of iron(III) oxide (rust) is produced from a certain amount of iron, how many
grams of oxygen are needed for this reaction?
A) 3.46 g
B) 1.54 g
C) 2.30 g
D) 6.91 g
E) none of these
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56. For the reaction
CaCO3 (s)  2HCl(aq)  CaCl2 (aq)  CO 2 (g)  H 2 O(l)
how many grams of CaCl2 can be obtained if 14.2 g HCl is allowed to react with excess
CaCO3?
A) 21.6 g CaCl2
B) 43 g CaCl2
C) 86 g CaCl2
D) 0.195 g CaCl2
E) none of these
57. True or false? The frequency of the wave is the distance between two consecutive wave
peaks.
A) True
B) False
58. The form of EMR that has less energy per photon than infrared rays but more energy per
photon than radio waves is
A) gamma rays
B) X rays
C) untraviolet
D) microwaves
E) none of these
59. Which color of visible light has the least amount of energy per photon?
A) green
B) violet
C) yellow
D) red
E) blue
60. As the principal energy level increases in an atom's orbitals, the average distance of an
electron energy level from the nucleus ______________.
A) varies
B) increases
C) decreases
D) stays the same
E) none of these
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61. A given set of f orbitals consists of ______________ orbital(s).
A) 9
B) 1
C) 3
D) 5
E) 7
62. The maximum electron capacity of an f sublevel is
A) 10
B) 18
C) 2
D) 14
E) 6
63. The maximum number of electrons allowed in the p sublevel of the third principal level
is
A) 6
B) 3
C) 1
D) 8
E) 2
64. The maximum number of electrons allowed in the fourth energy level is
A) 32
B) 8
C) 18
D) 2
E) 4
65. The number of d orbitals in the second principal energy level is
A) 14
B) 10
C) 2
D) 6
E) none of these
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66. The number of unpaired electrons in a nitrogen atom is
A) 4
B) 1
C) 2
D) 5
E) 3
67. Choose the correct ground-state electron configuration for oxygen.
A) [He]3s23p4
B) [He]1s22p6
C) [He]2s22p6
D) [Ne]2s22p4
E) [He]2s22p4
68. The alkaline earth metals have how many valence electrons?
A) 3
B) 8
C) 1
D) 7
E) 2
69. When moving down a group (family) in the periodic table, the number of valence
electrons
A) increases by 2 then 8 then 18 then 32
B) doubles with each move
C) changes in an unpredictable manner
D) remains constant
E) decreases regularly
70. Which of the following atoms has the electron configuration 1s22s22p63s23p64s23d1?
A) Sr
B) Sc
C) Ar
D) Ca
E) none of these
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71. Which electron configuration indicates a transition element?
A) 1s22s22p63s23p64s23d104p2
B) 1s22s22p63s13p6
C) 1s22s22p63s23p64s23d3
D) 1s22s22p5
E) none of these
72. How many of the following electron configurations for the species in their ground state
are correct?
I. Ca: 1s22s22p63s23p64s2
II. Mg: 1s22s22p63s1
III. V: [Ar] 3s23d3
IV. As: [Ar] 4s23d104p3
V. P: 1s22s22p63p5
A) 4
B) 3
C) 5
D) 2
E) 1
73. Write the electron configuration for Cd.
74. Write the electron configuration for Cl.
75. Which of the following is ranked in order of largest to smallest atomic radius?
A) Rb > Ge > Mn > F > S
B) Rb > Mn > S > Ge > F
C) Rb > Mn > Ge > S > F
D) Mn > Rb > F > S > Ge
E) F > S > Ge > Mn > Rb
76. Which of the following atoms has the highest ionization energy?
A) P
B) Sb
C) Si
D) Al
E) As
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77. True or false: Covalent bonding occurs when electrons are shared by nuclei.
A) True
B) False
78. True or false? The greater the difference in electronegativity between two bonded
atoms, the more polar the bond.
A) True
B) False
79. True or false? N2 is an example of a covalent bond.
A) True
B) False
80. The most electronegative element of those listed is
A) Cs
B) K
C) Li
D) Fr
E) Rb
81. Which of the following has nonpolar bonds?
A) HCl
B) Br2
C) All are nonpolar.
D) H2S
E) OF2
82. The number of polar covalent bonds in NH3 is
A) 3
B) 1
C) 4
D) 2
E) none of these
83. Draw the Lewis electron structure for the sulfur atom.
84. Draw the Lewis structure for SiH4.
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85. Draw the Lewis structure for CO.
86. Which of the following has a double bond?
A) CO
B) O2
C) H2S
D) NH3
E) H2O
Use the following to answer questions 87-88:
Use the following choices to describe the molecular structure of each of the following molecules
or ions.
a. linear
b. trigonal planar
c. tetrahedral
d. trigonal pyramid
e. Bent (V-shaped)
87. CH4
88. PF3
89. Convert 8.7  102 atm to torr.
A) 6.6  105 torr
B) 1.3  105 torr
C) 8.8  104 torr
D) 60 torr
E) 1.14 torr
90. The air in the inner tube of the tire of a racing bike has a pressure of 105.5 psi. Convert
this pressure to atm.
A) 1.041 atm
B) 105.5 atm
C) 0.1388 atm
D) 7.177 atm
E) 1551 atm
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91. Consider a gas at 1.00 atm in a 5.00-L container at 20.oC. What pressure does the gas
exert when transferred to a volume of 2.71 L at 43oC?
A) 0.585 atm
B) 0.371 atm
C) 1.99 atm
D) 3.97 atm
E) 1.71 atm
92. When analyzing ideal gases, the temperature must be measured on the Kelvin scale
A) because otherwise you could calculate a negative volume.
B) so that you are using an absolute scale.
C) to directly measure the average kinetic energy of the gas particles.
D) Both a and b are correct.
E) a, b, and c are correct.
93. Determine the pressure exerted by 4.01 mol of gas in a 2.92-L container at 32oC.
A) 10.53 atm
B) 34.4 atm
C) 30.8 atm
D) 100.4 atm
E) 3.61 atm
94. A sample of an ideal gas containing 0.779 mol is collected at 742 torr pressure and
31oC. Calculate the volume.
A) 2.67  10–3 L
B) 19.9 L
C) 2.03 L
D) 2.62  10–2 L
E) none of these
95. 3.31 mol of CO2 at STP will occupy
A) 37.1 L
B) 74.2 g
C) 80.9 L
D) 74.2 L
E) 9.76  10–2 L
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96. A vessel with an internal volume of 12.5 L contains 2.80 g of nitrogen gas, 0.403 g of
hydrogen gas, and 79.9 g of argon gas. At 25oC, what is the pressure (in atm) inside the
vessel?
A) 0.38 atm
B) 0.222 atm
C) 4.50 atm
D) 163 atm
E) 703 atm
97. What volume of HCl(g) measured at STP can be produced from 2.67 g of H2 and excess
Cl2 according to the following equation?
H 2 (g)  Cl2 (g)  2HCl(g)
A) 59.3 L
B) 120 L
C) 241 L
D) 89.7 L
E) 29.7 L
98. At 1 atm of pressure and a temperature of 0°C, which phase(s) of H2O can exist?
A) ice only
B) water vapor only
C) ice and water
D) water only
E) ice and water vapor
99. The normal boiling point of water is
A) 273 K
B) 373 K
C) 32°F
D) 0°F
E) none of these
100. True or false? The bonding forces that hold the atoms of a molecule together are called
intermolecular forces, whereas the forces that occur among molecules that cause them to
aggregate to form a solid or a liquid are called intramolecular forces.
A) True
B) False
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101. The bonds between hydrogen and oxygen in a water molecule can be characterized as
______________.
A) intramolecular forces
B) intermolecular forces
C) hydrogen bonds
D) dispersion forces
E) London forces
102. When a water molecule forms a hydrogen bond with another water molecule, which
atoms are involved in the interaction?
A) a hydrogen from one molecule and an oxygen from the other molecule
B) two hydrogens from one molecule and one hydrogen from the other molecule
C) two hydrogens from one molecule and one oxygen from the other molecule
D) a hydrogen from one molecule and a hydrogen from the other molecule
E) an oxygen from one molecule and an oxygen from the other molecule
103. Which of the following should have the lowest boiling point?
A) C3H8
B) C4H10
C) C5H12
D) CH4
E) C2H6
Use the following to answer questions 104-108:
Identify the major attractive force between particles in each of the compounds in questions as
one the following:
dipole-dipole
hydrogen bonding
London dispersion
ionic
104. O2
105. Octane (C8H18)
106. CH3CH2OH
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107. Sodium chloride
108. CO
109. Consider the following compounds:
CO
NH3
CO2
CH4
H2
How many of the compounds above exhibit London dispersion forces?
A) 2
B) 3
C) 5
D) 1
E) 4
110. Rank the following compounds from lowest to highest boiling point.
CH3OH
CH4
H2O
A) C2H6 < CH4 < CH3OH < H2O
B) CH4 < C2H6 < CH3OH < H2O
C) CH4 < C2H6 < H2O < CH3OH
D) CH4 < CH3OH < C2H6 < H2O
E) H2O < CH3OH < C2H6 < CH4
C2H6
111. In an atomic solid, what are the individual components?
A) atoms
B) electrons
C) bonds
D) molecules
E) ions
Use the following to answer questions 112-117:
Identify the type of solid as one of the following:
Ionic Solid
Molecular Solid
Atomic Solid
112. dry ice (solid carbon dioxide)
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113. xenon
114. copper metal
115. SF6
116. sulfur (S8)
117. NaHCO3
118. The total mass of a solution is 163.8 g. The solvent mass is 125.2 g. What is the mass
percent of the solute?
A) Not enough information is given.
B) 23.6 %
C) 76.4 %
D) 34.6 %
E) 15.8 %
119. A 120.0-g sample of nitric acid solution that is 70.0% HNO3 (by mass) contains
A) 5.29  103 mol HNO3
B) 84.0 mol HNO3
C) 1.33 mol HNO3
D) 1.90 mol HNO3
E) none of these
120. Determine the concentration of a solution made by dissolving 15.8 g of sodium chloride
in 750.0 mL of solution.
A) 0.360 M
B) 0.270 M
C) 11.9 M
D) 21.1 M
E) 0.203 M
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121. A solution is prepared by dissolving 5.57 g of Na2SO4 in enough water to make 225 mL
of solution. Calculate the solution molarity.
A) 0.174 M
B) 0.399 M
C) 1.25 M
D) 0.518 M
E) 0.0392 M
122. How many grams of CaCl2 (molar mass = 111.0 g/mol) are needed to prepare 4.97 L of
0.500 M CaCl2 solution?
A) 326 g
B) 301 g
C) 399 g
D) 288 g
E) 276 g
123. What mass of calcium chloride, CaCl2, is needed to prepare 3.995 L of a 1.56 M
solution?
A) 692 g
B) 284 g
C) 6.23 g
D) 471 g
E) 43.3 g
124. A 88.02-g sample of NaCl is dissolved in 250.0 mL of solution. Calculate the molarity
of this solution.
A) 352.1 M
B) 7.258 M
C) 1.506 M
D) 6.025 M
E) none of these
125. What volume of 12.0 M nitric acid is required to prepare 3.65 L of 0.100 M nitric acid?
A) 32.9 L
B) 3.29 L
C) 0.329 L
D) 0.365 L
E) 0.0304 L
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126. A 56.08-g sample of Ba(OH)2 is dissolved in enough water to make 1.20 L of solution.
How many milliliters of this solution must be diluted with water in order to make 1.00 L
of 0.100 M Ba(OH)2? (Ignore significant figures for this problem.)
A) 327 mL
B) 561 mL
C) 273 mL
D) 67.3 mL
E) 367 mL
127. Assume that vinegar is a 0.852 M solution of acetic acid (HC2H3O2) in water. What
volume of 0.2136 M NaOH would be needed to completely neutralize 5.92 mL of
vinegar?
A) 23.6 mL
B) 4.00 mL
C) 1.26 mL
D) 1.48 mL
E) 5.04 mL
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Answer Key
1. C
Chapter: Ch 2.5
2. D
Chapter: Ch 2.2,2.6
3. A
Chapter: Ch 2.2,2.6
4. A
Chapter: Ch 2.7
5. D
Chapter: Ch 2.7
6. B
Chapter: Ch 2.7
7. D
Chapter: Ch 2.8
8. C
Chapter: Ch 2.6
9. C
Chapter: Ch 3.2
10. A
Chapter: Ch 3.3
11. D
Chapter: Ch 3.4
12. The isotope contains 26 protons, 23 electrons, and 31 neutrons.
Chapter: Ch 4.7,4.10
13. A
Chapter: Ch 4.6
14. C
Chapter: Ch 4.8
15. E
Chapter: Ch 4.4
16. E
Chapter: Ch 5.6
17. Ca(HCO3)2
Chapter: Ch 5.5
18. N2O5
Chapter: Ch 5.3
19. nitrate ion
Chapter: Ch 5.5
20. CO
Chapter: Ch 5.3
21. HClO
Chapter: Ch 5.6
22. ammonium sulfate
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23.
24.
25.
26.
27.
28.
29.
30.
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
41.
42.
43.
44.
45.
Chapter: Ch 5.5
acetate ion
Chapter: Ch 5.5
H2S
Chapter: Ch 5.6
ammonium carbonate
Chapter: Ch 5.5
CrI3
Chapter: Ch 5.2
zinc hydroxide
Chapter: Ch 5.5
A
Chapter: Ch 5.3
D
Chapter: Ch 6.3
2H2O2(l)  2H2O(l) + O2(g)
Chapter: Ch 6.3
b
Chapter: Ch 7.6
b
Chapter: Ch 7.6
b; c
Chapter: Ch 7.6
a
Chapter: Ch 7.7
a; c
Chapter: Ch 7.7
a; b
Chapter: Ch 7.7
a; d
Chapter: Ch 7.7
A
Chapter: Ch 8.7
D
Chapter: Ch 8.3
C
Chapter: Ch 8.3
B
Chapter: Ch 8.3
C
Chapter: Ch 8.3
D
Chapter: Ch 8.5
D
Chapter: Ch 8.3
D
Page 24
46.
47.
48.
49.
50.
51.
52.
53.
54.
55.
56.
57.
58.
59.
60.
61.
62.
63.
64.
65.
66.
67.
68.
Chapter:
C
Chapter:
D
Chapter:
C
Chapter:
E
Chapter:
D
Chapter:
E
Chapter:
C2H3O2
Chapter:
B
Chapter:
D
Chapter:
A
Chapter:
A
Chapter:
B
Chapter:
D
Chapter:
D
Chapter:
B
Chapter:
E
Chapter:
D
Chapter:
A
Chapter:
A
Chapter:
E
Chapter:
E
Chapter:
E
Chapter:
E
Ch 8.3
Ch 8.5
Ch 8.5
Ch 8.8
Ch 8.6
Ch 8.7
Ch 8.8
Ch 8.8
Ch 9.2
Ch 6.3,9.3
Ch 9.3
Ch 9.3
Ch 11.2
Ch 11.2
Ch 11.2
Ch 11.4
Ch 11.8
Ch 11.8
Ch 11.8
Ch 11.8
Ch 11.8
Ch 11.9
Ch 11.9
Page 25
69.
70.
71.
72.
73.
74.
75.
76.
77.
78.
79.
80.
81.
82.
Chapter: Ch 11.10
D
Chapter: Ch 11.11
B
Chapter: Ch 11.10
C
Chapter: Ch 11.10
D
Chapter: Ch 11.10
[Kr] 5s24d10
Chapter: Ch 11.10
[Ne] 3s23p5
Chapter: Ch 11.10
C
Chapter: Ch 11.11
A
Chapter: Ch 11.11
A
Chapter: Ch 12.1
A
Chapter: Ch 12.2
A
Chapter: Ch 12.1
C
Chapter: Ch 12.2
B
Chapter: Ch 12.3
A
Chapter: Ch 12.3
83. . . .
. S.
.
Chapter: Ch 12.6
84.
H
—
H — Si — H
—
H
Chapter: Ch 12.6
.
85. .. C —
—
— O.
Chapter: Ch 12.7
86. B
Chapter: Ch 12.7
87. c
Page 26
88.
89.
90.
91.
92.
93.
94.
95.
96.
97.
98.
99.
100.
101.
102.
103.
104.
105.
106.
107.
108.
109.
110.
111.
Chapter: Ch 12.9
d
Chapter: Ch 12.9
A
Chapter: Ch 13.1
D
Chapter: Ch 13.1
C
Chapter: Ch 13.5
E
Chapter: Ch 13.3
B
Chapter: Ch 13.5
B
Chapter: Ch 13.5
D
Chapter: Ch 13.5
C
Chapter: Ch 13.6
A
Chapter: Ch 13.10
C
Chapter: Ch 14.1
B
Chapter: Ch 14.1
B
Chapter: Ch 14.3
A
Chapter: Ch 14.3
A
Chapter: Ch 14.3
D
Chapter: Ch 14.3
London dispersion
Chapter: Ch 14.3
London dispersion
hydrogen bonding
ionic
dipole-dipole
Chapter: Ch 14.3
C
Chapter: Ch 14.3
B
Chapter: Ch 14.3
A
Chapter: Ch 14.5,14.6
Page 27
112.
113.
114.
115.
116.
117.
118.
119.
120.
121.
122.
123.
124.
125.
126.
127.
molecular solid
atomic solid
atomic solid
molecular solid
molecular solid
ionic solid
B
Chapter: Ch 15.3
C
Chapter: Ch 15.3
A
Chapter: Ch 15.4
A
Chapter: Ch 15.4
E
Chapter: Ch 15.4
A
Chapter: Ch 15.4
D
Chapter: Ch 15.4
E
Chapter: Ch 15.5
E
Chapter: Ch 15.5
A
Chapter: Ch 15.7
Page 28