Chemistry Final Exam Review MC
PSI Chemistry
Name___________________________
Atoms, Molecules and Ions
1. When a metal and a nonmetal react, the _____ tends to lose electrons and the _____ tends to gain
electrons.
A) metal, metal
B) nonmetal, nonmetal
C) metal, nonmetal
D) nonmetal, metal
E) None of the above, these elements share electrons.
2. What is the formula of the compound formed between strontium ions and nitrogen ions?
A) SrN
B) Sr3N2
C) Sr2N3
D) SrN2
E) SrN3
3. Barium forms an ion with a charge of __________.
A) 1+
B) 2D) 3E) 2+
4.
C) 3+
Predict the empirical formula of the ionic compound that forms from aluminum and oxygen.
A) AlO
B) Al3O2
C) Al2O3
D) AlO2
E) Al2O
5. The correct name for SrO is __________.
A) strontium oxide
B) strontium hydroxide
D) strontium monoxide
E) strontium dioxide
6. The name of PCl3 is __________.
A) potassium chloride
chloride
D) monophosphorous trichloride
B) phosphorus trichloride
C) strontium peroxide
C) phosphorous(III)
E) trichloro potassium
7. The ions Ca2+ and PO43- form a salt with the formula __________.
A) CaPO4
B) Ca2(PO4)3
D) Ca(PO4)2
E) Ca3(PO4)2
C) Ca2PO4
8. The formula of ammonium carbonate is __________.
A) (NH4)2CO3
B) NH4CO2
D) (NH3)2CO3
E) N2(CO3)3
C) (NH3)2CO4
9. Cathode rays are __________.
A) neutrons
D) protons
C) electrons
B) x-rays
E) atoms
10. There are __________ electrons, __________ protons, and __________ neutrons in an atom of
A) 132, 132, 54
B) 54, 54, 132
C) 78, 78, 54
D) 54, 54, 78
E) 78, 78, 132
132
54𝑋𝑒 .
11. Isotopes are atoms that have the same number of _______ but differing number of _______.
A) protons, electrons
B) neutrons, protons
C) protons, neutrons
D) electrons, protons
E) neutrons, electrons
12. The element X has three naturally occurring isotopes. The isotopic masses
(amu) and % abundances of the isotopes are given in this table. The
average atomic mass of the element is __________ amu.
A) 33.33
B) 55.74
C) 56.11
D) 57.23
E) 56.29
13. Of the following, __________ contains the greatest number of electrons.
A) P3+
B) P
D) P3E) P2+
C) P2-
14. Which of the following compounds would you expect to be ionic?
A) SF6
B) H2O
D) NH3
E) CaO
C) H2O2
15. Which one of the following compounds is chromium(III) oxide?
A) Cr2O3
B) CrO3
D) Cr3O
E) Cr2O4
C) Cr3O2
Structure of Atoms
16. What is the frequency (Hz) of electromagnetic radiation that has a wavelength of 0.53 m?
A) 5.7 x 108
B) 1.8 x 10-9
C) 1.6 x 108
-33
33
D) 1.3 x 10
E) 1.3 x 10
17. The energy of a photon that has a wavelength of 12.3 nm is __________ J.
A) 1.51 x 10-17
B) 4.42 x 10-23
-50
D) 2.72 x 10
E) 1.62 x 10-17
C) 1.99 x 10-25
18. Of the following, __________radiation has the shortest wavelength and __________radiation has the
greatest energy: gamma
ultraviolet
visible
A) gamma, visible
B) visible, gamma
C) visible, ultraviolet
D) ultraviolet, gamma
E) gamma, gamma
19. Which one of the following is the correct electron configuration for a ground-state nitrogen atom?
A)
B)
C)
D)
E) None of the above is correct.
20. What is the de Broglie wavelength (m) of a 2.0 kg object moving at a speed of 50 m/s?
A) 6.6 x 10-36
B) 1.5 x 1035
C) 5.3 x 10-33
D) 2.6 x 10-35
E) 3.8 x 1034
21. The __________ quantum number defines the shape of an orbital.
A) spin
B) magnetic
D) magnetic
E) phi
C) principal
22. There are __________ orbitals in the third shell.
A) 25
B) 4
D) 16
E) 1
C) 9
23. The n = 1 shell contains ______p orbitals. All the other shells contain _____ p orbitals.
A) 3, 6
B) 0, 3
C) 6, 2
D) 3, 3
E) 0, 6
24. The lowest energy shell that contains f orbitals is the shell with n = __________.
A) 3
B) 2
C) 4
D) 1
E) 5
25. How many p-orbitals are occupied in a Ne atom?
A) 5
B) 6
D) 3
E) 2
C) 1
26. How many quantum numbers are necessary to designate a particular electron in an atom?
A) 3
B) 4
C) 2
D) 1
E) 5
27. There are __________ unpaired electrons in a ground state phosphorus atom.
A) 0
B) 1
D) 3
E) 4
C) 2
28. The ground state electron configuration for Zn is __________.
A) [Kr] 4s2 3d10
B) [Ar] 4s2 3d10
D) [Ar] 3s2 3d10
E) [Kr] 3s2 3d10
C) [Ar] 4s1 3d10
29. At maximum, an f-subshell can hold __________ electrons, a d-subshell can hold __________ electrons,
and a p-subshell can hold __________ electrons.
A) 14, 10, 6
B) 2, 8, 18
C) 14, 8, 2
D) 2, 12, 21
E) 2, 6, 10
30. Which one of the following represents an acceptable set of quantum numbers for an electron in an atom?
(arranged as n, l, ml, and ms)
A) 2, 2, -1, -1/2
B) 1, 0, 0, 1/2
C) 3, 3, 3, 1/2
D) 5, 4, -5, 1/2
E) 3, 3, 3, -1/2
Periodic Properties
31. The first ionization energies of the elements __________ as you go from left to right across a period of
the periodic table, and __________ as you go from the bottom to the top of a group in the table.
A) increase, increase
B) increase, decrease
C) decrease, increase
D) decrease, decrease
E) are completely unpredictable
32. The most common sulfur ion has a charge of __________.
A) 2B) 1D) 6+
E) Sulfur does not form ions.
C) 4+
33. In which set of elements would all members be expected to have very similar chemical properties?
A) O, S, Se
B) N, O, F
C) Na, Mg, K
D) S, Se, Si
E) Ne, Na, Mg
34. Of the following, which gives the correct order for atomic radius for Mg, Na, P, Si and Ar?
A) Mg > Na > P > Si > Ar
B) Ar > Si > P > Na > Mg
C) Si > P > Ar > Na
> Mg
D) Na > Mg > Si > P > Ar
E) Ar > P > Si > Mg > Na
35. The atomic radius of main-group elements generally increases down a group because __________.
A) effective nuclear charge increases down a group
B) effective nuclear charge decreases down a group
C) effective nuclear charge zigzags down a group
D) the principal quantum number of the valence orbitals increases
E) both effective nuclear charge increases down a group and the principal quantum number of the valence
orbitals increases
36. Which one of the following atoms has the largest radius?
A) I
B) Co
D) Sr
E) Ca
C) Ba
37. In which of the following atoms is the 2s orbital closest to the nucleus?
A) S
B) Cl
D) Si
E) It’s the same in all of these atoms.
C) P
38. Which of the following is an isoelectronic series?
A) B5-, Sr4-,As3-, Te2B) F-, Cl-, Br-, ID) Si2-, P2-, S2-, Cl2E) O2-, F-, Ne, Na+
C) S, Cl, Ar, K
39. Which of the following correctly lists the atoms in order of increasing size (smallest to largest)?
A) F < K < Ge < Br < Rb
B) F < Ge < Br < K < Rb
C) F < K < Br < Ge
< Rb
D) F < Br < Ge < K < Rb
E) F < Br < Ge < Rb < K
40. Of the choices below, which gives the order for first ionization energies?
A) Cl > S > Al > Ar > Si
B) Ar > Cl > S > Si > Al
> Ar
D) Cl > S > Al > Si > Ar
E) S > Si > Cl > Al > Ar
C) Al > Si > S > Cl
41. Of the following elements, which has the largest first ionization energy?
A) Se
B) As
D) Sb
E) Ge
C) S
42. Which of the following has the largest second ionization energy?
A) Si
B) Mg
D) Na
E) P
C) Al
43. Which ion below has the largest radius?
A) ClB) K+
D) F
E) Na+
C) Br-
44. Of the following elements, __________ has the greatest electronegativity.
A) S
B) Cl
D) Br
E) I
C) Se
45. Of the elements below, __________ is the most metallic.
A) Na
B) Mg
D) K
E) Ar
C) Al
Chemical Bonding
46. The halogens, alkali metals, and alkaline earth metals have __________ valence electrons, respectively.
A) 7, 4, and 6
B) 1, 5, and 7
C) 8, 2, and 3
D) 7, 1, and 2
E) 2, 7, and 4
47. What is the electron configuration for the Co2+ ion?
A) [Ar] 4s1 3d6
B) [Ar] 4s0 3d7
D) [Ar] 4s2 3d9
E) [Ne] 3s2 3p10
C) [Ar] 4s0 3d5
48. Elements from opposite sides of the periodic table tend to form __________.
A) covalent compounds
B) ionic compounds
C) compounds that are gaseous at room temperature
D) homonuclear diatomic compounds
E) covalent compounds that are gaseous at room temperature
49. A double bond consists of __________ pairs of electrons shared between two atoms.
A) 1
B) 2
C) 3
D) 4
E) 6
50. The Lewis structure of
A) 0
D) 3
AsH 3 shows __________ nonbonding electron pair(s) on As.
B) 1
C) 2
E) This cannot be determined from the data given.
51. How many different types of resonance structures can be drawn for the ion SO32- where all atoms satisfy
the octet rule?
A) 1
B) 2
C) 3
D) 4
E) 5
52. Which of the following does not have eight valence electrons?
A) Ca+
B) Rb+
C) Xe
D) Br
E) All of the above have eight valence electrons.
53. In the Lewis symbol for a sulfur atom, there are __ paired and __ unpaired electrons.
A) 2, 2
B) 4, 2
C) 2, 4
D) 0, 6
E) 5, 1
54. A valid Lewis structure of _______ cannot be drawn without violating the octet rule.
A) PO43B) SiF4
C) CF4
D) SeF4
E) NF3
Molecular Geometry
55. The molecular geometry of __________ is square planar.
A) CCl4
B) XeF4
D) XeF2
E) ICl3
56. The molecular geometry of the CS2 molecule is __________.
A) linear
B) bent
D) trigonal planar
E) T-shaped
C) PH3
C) tetrahedral
57. The electron domain and molecular geometry of BrO2- is _________.
A) tetrahedral, trigonal planar
B) trigonal planar, trigonal planar
linear
D) tetrahedral, bent
E) trigonal pyramidal, seesaw
58. The electron-domain geometry of __________ is tetrahedral.
A) CBr4
B) PH3
D) XeF4
E) all of the above except XeF4
C) trigonal pyramidal,
C) CCl2Br2
59. The central Xe atom in the XeF4 molecule has __________ unbonded electron pairs and __________
bonded electron pairs in its valence shell.
A) 1, 4
B) 2, 4
C) 4, 0
D) 4, 1
E) 4, 2
60. Of the molecules below, only __________ is nonpolar.
A) BF3
B) NF3
D) PBr3
E) BrCl3
C) IF3
61. Of the following molecules, only __________ is polar.
A) BeCl2
B) BF3
D) SiH2Cl2
E) Cl2
C) CBr4
62. Using the VSEPR model, the electron-domain geometry of the central atom in BrF4- is ______.
A) linear
B) trigonal planar
C) tetrahedral
D) trigonal bipyramidal
E) octahedral
63. Using the VSEPR model, the molecular geometry of the central atom in NCl3 is __________.
A) linear
B) trigonal planar
C) tetrahedral
D) bent
E) trigonal pyramidal
Stoichiometry
64. When the following equation is balanced, the coefficients are __________.
NH3(g) + O2 → NO2(g) + H2O(g)
A) 1, 1, 1, 1
B) 4, 7, 4, 6
D) 1, 3, 1, 2
E) 4, 3, 4, 3
C) 2, 3, 2, 3
65. When the following equation is balanced, the coefficient of HNO3 is __________.
HNO3(aq) + CaCO3(s) → Ca(NO3)2(aq) + CO2(g) + H2O(l)
A) 1
B) 2
C) 3
D) 5
E) 4
66. When methanol, CH3OH(l) is burned in air, what is the coefficient of methanol in the balanced equation?
A) 1
B) 2
C) 3
D) 4
E) 3/2
67. How many grams of oxygen are in 65 g of C2H2O2?
A) 18
B) 29
D) 36
E) 130
C) 9.0
68. A compound contains 38.7% K, 13.9% N, and 47.4% O by mass. What is the empirical formula of the
compound?
A) KNO3
B) K2N2O3
C) KNO2
D) K2NO3
E) K4NO5
69. Of the reactions below, which one is a decomposition reaction?
A) NH4Cl → NH3 + HCl
B) 2Mg + O2 → 2MgO
D) 2CH4 + 4O2 → 2CO2 + 4H2O
E) Cd(NO3)2 + Na2S →CdS + 2NaNO3
C) 2N2 + 3H2 → 2NH3
70. Which one of the following substances is the product of this combination reaction?
Al(s) + I2(s) → _____
A) AlI2
B) AlI
C) AlI3
D) Al2I3
E) Al2I3
71. Which of the following are combination reactions?
1) CH4(g) + O2(g) → CO2(g) + H2O(l)
2) CaO(s) + CO2(g) → CaCO3(s)
3) Mg(s) + O2(g) → MgO(s)
4) PbCO3(s) → PbO(s) + CO2(g)
A) 1, 2, and 3
B) 2 and 3
D) 4 only
E) 2, 3, and 4
C) 1, 2, 3, and 4
72. The formula of nitrobenzene is C6H5NO2. The molecular weight of this compound is _____ amu.
A) 107.11
B) 43.03
C) 109.10
D) 123.11
E) 3.06
73. The mass % of F in the binary compound KrF2 is __________.
A) 18.48
B) 45.38
D) 81.52
E) 31.20
C) 68.80
74. A 30.5 gram sample of glucose (C6H12O6) contains __________ mol of glucose.
A) 0.424
B) 0.169
C) 5.90
D) 2.36
E) 0.136
75. Sulfur and oxygen react to produce sulfur trioxide. In a particular experiment, 7.9 grams of SO3 are
produced by the reaction of 5.0 grams of O2 with 6.0 grams of S. What is the % yield of SO3 in this
experiment?
S(s) + O2(g) → SO3(g) (not balanced)
A) 32
B) 63
C) 75
D) 95
E) 99
76. The combustion of ammonia in the presence of excess oxygen yields NO2 and H2O:
4 NH3(g) + 7 O2(g) → 4 NO2(g) + 6 H2O(g)
The combustion of 43.9 g of ammonia produces __________ g of NO2.
A) 2.58
B) 178
C) 119
D) 0.954
E) 43.9
77. The combustion of propane (C3H8) produces CO2and H2O:
C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(g)
The reaction of 2.5 mol of O2 will produce __________ mol of H2O.
A) 4.0
B) 3.0
D) 2.0
E) 1.0
C) 2.5
78. What is the maximum amount in grams of SO3that can be produced by the reaction of 1.0 g of S with 1.0
g of O2 via the equation below?
S(s) + O2(g) → SO3(g)
(not balanced)
A) 0.27
B) 1.7
C) 2.5
D) 3.8
E) 2.0
79. Solid aluminum and gaseous oxygen react in a combination reaction to produce aluminum oxide:
4 Al(s) + 3 O2(g) → 2 Al2O3(s)
The maximum amount of Al2O3 that can be produced from 2.5 g of Al and 2.5 g of O2 is ________ g.
A) 9.4
B) 7.4
C) 4.7
D) 5.3
E) 5.0
Gases
80. A temperature of 373 K is equivalent to?
A) -100 oC
B) 0 oC
o
D)273 C
E) 373 oC
C) 100 oC
81. Which of the following equations shows an incorrect relationship between pressures given in terms of
different units?
A) 1.20 atm = 122 kPa
B) 152 mm Hg = 2.03 x 104
C) 0.760 atm = 578 mm
Hg
D) 1.0 torr = 2.00 mm Hg
E) 1.00 atm = 760 torr
82. At a temperature of __________ °C, 0.444 mol of CO gas occupies 11.8 L at 889 torr.
A) 379
B) 73
C) 14
D) 32
E) 106
83. How many moles of gas are there in a 45.0 L container at 25.0 °C and 500.0 mm Hg?
A) 0.630
B) 6.11
C) 18.4
D) 1.21
E) 207
84. A sample of a gas (5.0 mol) at 1.0 atm is expanded at constant temperature from 10 L to 15 L. The final
pressure is __________ atm.
A) 1.5
B) 7.5
C) 0.67
D) 3.3
E) 15
85. A gas originally at 27 °C and 1.00 atm pressure in a 3.9 L flask is cooled at constant pressure until the
temperature is 11 °C. The new volume of the gas is __________ L.
A) 0.27
B) 3.7
C) 3.9
D) 4.1
E) 0.24
86. A sample of H2 gas (12.28 g) occupies 100.0 L at 400.0 K and 2.00 atm. A sample weighing 9.49 g
occupies __________ L at 353 K and 2.00 atm.
A) 109
B) 68.2
C) 54.7
D) 147
E) 77.3
87. The density of N2O at 1.53 atm and 45.2 °C is __________ g/L.
A) 18.2
B) 1.76
D) 9.99
E) 2.58
C) 0.388
88. The density of nitric oxide (NO) gas at 1.21 atm and 54.1 °C is __________ g/L.
A) 0.0451
B) 0.740
C) 1.35
D) 0.273
E) 8.2
89. A mixture of He and Ne at a total pressure of 0.95 atm is found to contain 0.32 mol of He and 0.56 mol of
Ne. The partial pressure of Ne is __________ atm.
A) 1.7
B) 1.5
C) 0.60
D) 0.35
E) 1.0
90. Arrange the following gases in order of increasing average molecular speed at 25 °C.
Cl2, O2, F2, N2
A) Cl2 < F2 < O2 < N2
B) Cl2 < O2 < F2 < N2
C) N2 < F2 < Cl2 <
O2
D) Cl2 < F2 < N2 < O2
E) F2 < O2 < N2 < Cl2
91. A flask contains a mixture of He and Ne at a total pressure of 2.6 atm. There are 2.0 mol of He and 5.0
mol of Ne in the flask. The partial pressure of He is __________ atm.
A) 9.1
B) 6.5
C) 1.04
D) 0.74
E) 1.86
92. A sample of oxygen gas was found to effuse at a rate equal to two times that of an unknown gas. The
molecular weight of the unknown gas is __________ g/mol.
A) 64
B) 128
C) 8
D) 16
E) 8.0
93. Which noble gas is expected to show the largest deviations from the ideal gas behavior?
A) helium
B) neon
C) argon
D) krypton
E) xenon
Intermolecular Forces, Liquids and Solids
94. The strongest interparticle attractions exist between particles of a __________ and the weakest
interparticle attractions exist between particles of a __________.
A) solid, liquid
B) solid, gas
C) liquid, gas
D) liquid, solid
E) gas, solid
95. Of the following substances, only __________ has London dispersion forces as the only intermolecular
force.
A) CH3OH
B) NH3
C) H2S
D) Kr
E) HCl
96. The predominant intermolecular force in (CH3)2NH is __________.
A) London dispersion forces
B) ion-dipole forces
D) dipole-dipole forces
E) hydrogen bonding
C) ionic bonding
97. Which one of the following should have the lowest boiling point?
A) PH3
B) H2S
D) SiH4
E) H2O
C) HCl
98. What types of intermolecular forces exist between HI and H2S?
A) dipole-dipole and ion-dipole
B) dispersion forces, dipole-dipole, and ion-dipole
C) dispersion forces, hydrogen bonding, dipole-dipole, and ion-dipole
D) dispersion forces and dipole-dipole
E) dispersion forces, dipole-dipole, and ion-dipole
99. The substance with the largest heat of vaporization is __________.
A) I2
B) Br2
D) F2
E) O2
100.
Of the following, __________ is an exothermic process.
A) melting
B) subliming
D) boiling
E) All are exothermic.
C) Cl2
C) freezing
101.
A)
B)
C)
D)
Water could be made to boil at 105°C instead of 100°C by _____.
increasing the air pressure on the water
decreasing the air pressure above the water
decreasing the pressure on the water
applying a great deal of heat
102. On the phase diagram shown to the right, segment __________
corresponds to the conditions of temperature and pressure under which
the solid and the gas of the substance are in equilibrium.
A) AB B) AC C) AD D) CD E) BC
103. On the phase diagram shown to the right, the coordinates of point
________ correspond to the critical temperature and pressure.
A) A B) B C) C D) D E) E
104. How much heat does it take to warm 16.0 g of pure water from 90.0°C to 100.0°C? (specific heat of water
= 4.18 J/g x °C)
A) 66.9 joules
B) 669 joules
C)16.0 joules
D) 160 joules
105. 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 (cal/g x °C))
A)3.8 kcal
B) 85 kcal
C)0.59 kcal
D) none of the above
Thermochemistry
106. The value of ΔE for a system that performs 13 kJ of work on its surroundings and loses 9 kJ of heat is
A) 22 kJ
B) -22 kJ
C) -4 kJ
D) 4 kJ
E) -13 kJ
107. When a system __________, ΔE is always negative.
A) absorbs heat and does work
B) gives off heat and does work
C) absorbs heat and has work done on it
D) gives off heat and has work done on it
E) none of the above
108. ΔH for an endothermic process is _____ while ΔH for an exothermic process is _______.
A) zero, positive
B) zero, negative
C) positive, zero
D) negative, positive
E) positive, negative
109. An 8.29 g sample of calcium carbonate [CaCO3(s)] absorbs 50.3 J of heat, upon which the temperature of
the sample increases from 21.1 °C to 28.5 °C. What is the specific heat of calcium carbonate?
A) 0.63
B) 0.82
C) 1.1
D) 2.2
E) 4.2
110. A substance releases 500 kJ of heat as 25 mol of it condenses from a gas to a liquid . What is the heat of
vaporization (ΔHvap) of this substance?
A) 20 kJ/mol
B) 25 kJ/mol
C) 475 kJ/mol
D) 525 kJ/mol
111. The heat of fusion of water is 6.01 kJ/mol. The heat capacity of liquid water is 75.3 J/mol-K The
conversion of 50.0 g of ice at 0.00 °C to liquid water at 22.0 °C requires __________ kJ of heat.
A) 3.8 x 102
B) 21.3
C) 17.2
D) 0.469
E) Insufficient data are given.
112. The value of ΔHo for the reaction CH4(g) + 3Cl2(g)  CHCl3(l) + 3HCl(g) is -336 kJ. Calculate the heat (kJ)
released to the surroundings when 23.0 g of HCl is formed.
A) 177
B) 2.57 x 103
C) 70.7
D) 211
E) -336
113. The value of ΔHo for the reaction H2(g) + Cl2  2HCl(g) is -186 kJ. Calculate the heat (kJ) released from
the reaction of 25 g of Cl2.
A) 66
B) 5.3 x 102
C) 33
D) 47
E) -186
114. ΔH for the reaction IF5(g)  IF3(g) + F2(g) is __________ kJ, give the data below.
IF(g) + F2(g)  IF3(g)
ΔH = -390 kJ
IF(g) + 2F2(g)  IF5(g)
ΔH = -745 kJ
A) +355
B) -1135
D) +35
E) -35
C) +1135
115. ΔH for the reaction 3Fe2O3(s) + CO(g)  CO2(g) + 2Fe3O4(s) is __________ kJ, give the data below.
Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g) ΔH = -28.0 kJ
3Fe(s) + 4CO2(g)  4CO(g) + Fe3O4(s)
ΔH = +12.5 kJ
A) -59.0
B) 40.5
C) -15.5
D) -109
E) +109
116. The value of ΔHo for the reaction H2(g) + Cl2(g)  2HCl(g) is -186 kJ. The value of ΔHfo for HCl (g) is
__________ kJ/mol.
A) -3.72 x 102
B) -1.27 x 102
C) -93.0
D) -186
E) +186
117. The value of ΔHo for the reaction 2Al(s) +3O2(g)  2Al2O3(s) is -3351 kJ. The value of ΔHfo for Al2O3(s) is
A) -3351 kJ.
B) -1676 kJ.
C) -32.86 kJ.
D) -16.43 kJ.
E) +3351 kJ.
Thermodynamics
118. A reaction that is spontaneous __________.
A) is very rapid
B) will proceed without outside intervention
C) is also spontaneous in the reverse direction
D) has an equilibrium position that lies far to the left
E) is very slow
119. Of the following, the entropy of gaseous __________ is the largest at 25 oC and 1 atm.
A) H2 B) C2H6
C) C2H2
D) CH4
E) C2H4
120. ΔS is positive for the reaction __________.
A) 2H2(g) + O2(g) → 2H2O(g)
B) 2NO(g) → N2O4(g)
D) BaF2(s) → Ba2+ + 2F-(aq)
E) 2Hg(l) + O2(g) → 2HgO(s)
C) CO2(g) → CO2(s)
121. ΔS is negative for the reaction __________.
A) 2SO2(g) + O2(g) → 2SO3(g)
B) NH4Cl(s) → NH3(g) + HCl(g)
2ClD) 2C(s) + 2O2(g) → 2CO2(g)
E) H2O(l) → H2O(g)
C) PbCl2(s) → Pb2+ +
122. The value of ΔSo for the formation of POCl3 from its constituent elements, as shown below, is
P2(g) + O2(g) + 3Cl2(g)  2POCl3(g)
A) -442.0 J/K∙ mol.
B) +771.0 J/K∙ mol.
C) -321.0 J/K∙ mol.
D) -771.0 J/K∙ mol.
E) +321.0 J/K∙ mol.
123. The value of ΔSo for the formation of calcium chloride from its constituent elements, as shown below, is
Ca(s) + Cl2(g)  CaCl2(s)
A) -104.6 J/K∙ mol.
B) +104.6 J/K∙ mol.
C) +369.0 J/K∙ mol.
D) -159.8 J/K∙ mol.
E) +159.8 J/K∙ mol.
124. The value of ΔGo at 25oC for the formation of POCl3 from its constituent elements, as shown below, is
P2(g) + O2(g) + 3Cl2(g)  2POCl3(g)
A) -1,108.7 kJ/mol.
B) +1,108.7 kJ/mol.
C) -606.2 kJ/mol.
D) +606.2 kJ/mol.
E) -1,005 kJ/mol.
125. The value of ΔGo at 25oC for the formation of PCl3from its constituent elements, as shown below, is
P2(g) + 3Cl2(g)  2PCl3(g)
A) -539.2 kJ/mol.
B) +539.2 kJ/mol.
C) -642.9 kJ/mol.
D) +642.9 kJ/mol.
E) -373.3 kJ/mol.
126. A reaction that is not spontaneous at low temperature can become spontaneous at high temperature if
ΔH is __________ and ΔS is __________.
A) +, +
B) -, C) +, D) -, +
E) +, 0
127. For the below reaction, ΔHo = 131.3 kJ/mol and ΔSo = 133.6 J/mol at 298 K. At temperatures greater
than _______°C this reaction is spontaneous under standard conditions
C(s) + H2O(g) CO(g) + H2(g)
A) 273
B) 325
C) 552
D) 710
E) 983
128. Given the following table of thermodynamic data, the vaporization of TiCl4 is __________.
A)
B)
C)
D)
E)
spontaneous at all temperatures
spontaneous at low temperature and nonspontaneous at high temperature
nonspontaneous at low temperature and spontaneous at high temperature
nonspontaneous at all temperatures
not enough information given to draw a conclusion
Solutions
129. A small amount of salt dissolved in water is an example of a __________.
A) homogeneous mixture
B) heterogeneous mixture
D) pure substance
E) solid
C) compound
130. Which one of the following is a pure substance?
A) concrete
B) wood
D) elemental copper
E) milk
C) salt water
131. Homogeneous mixtures are also known as __________.
A) solids
B) compounds
D) substances
E) solutions
C) elements
132. Of the following, only __________ is an extensive property.
A) density
B) mass
D) freezing point
E) temperature
C) boiling point
133. Mass and volume are often referred to as __________ properties of substances.
A) extensive
B) intensive
C) chemical
D) heterogeneous
134. A strong electrolyte is one that __________ completely in solution.
A) reacts
B) associates
C) disappears
D) ionizes
135. A weak electrolyte exists predominantly as __________ in solution.
A) atoms
B) ions
D) electrons
E) an isotope
C) molecules
136. The process of solute particles being surrounded by solvent particles is known as _____.
A) salutation
B) agglomeration
C) solvation
D) agglutination
E) dehydration
137. A saturated solution __________.
A) contains as much solvent as it can hold
B) contains no double bonds
C) contains dissolved solute in equilibrium with undissolved solute
D) will rapidly precipitate if a seed crystal is added
E) cannot be attained
138. A supersaturated solution __________.
A) is one with more than one solute
B) is one that has been heated
C) is one with a higher concentration than the solubility
D) must be in contact with undissolved solid
E) exists only in theory and cannot actually be prepared
139. Which one of the following substances would be the most soluble in CCl4?
A) CH3CH2OH
B) H2O
D) C10H22
E) NaCl
C) NH3
140. Which of the following substances is more likely to dissolve in CH 3OH ?
A) CCl4
B) Kr
D) CH3CH2OH
E) H2
C) N2
141. Which one of the following is most soluble in water?
A) CH3OH
B) CH3CH2OH
D) CH3CH2CH2CH2OH
E) CH3CH2CH2CH2CH2OH
C) CH3 CH2CH2OH
142. The concentration of urea in a solution prepared by dissolving 16 g of urea in 39 g of H 2O is
__________% by mass.
A) 29
B) 41
C) 0.29
D) 0.41
E) 0.48
143. A solution is prepared by dissolving 15.0 g of NH3in 250.0 g of water. The mole fraction of NH3in the
solution is __________.
A) 0.0640
B) 0.0597
C) 0.940
D) 0.922
E) 16.8
144. When 0.500 mol of HC2H3O2 is combined with enough water to make a 300.0 mL solution, the
concentration of HC2H3O2 is __________ M.
A) 3.33
B) 1.67
C) 0.835
D) 0.00167
E) 0.150
145. What is the concentration (M) of a NaCl solution prepared by dissolving 9.3 g of NaCl in sufficient water to
give 350 mL of solution?
A) 18
B) 0.16
C) 0.45
D) 27
E) 2.7 x 10-2
146. The molarity (M) of an aqueous solution containing 52.5 g of sucrose (C 12H22O11) in 35.5 mL of solution is
A) 5.46
B) 1.48
C) 0.104
D) 4.32
E) 1.85
147. The concentration of KBr in a solution prepared by dissolving 2.21 g of KBr in 897 g of water is ____
molal.
A) 2.46
B) 0.0167
C) 0.0207
D) 2.07 x 10-5
E) 0.0186
148. The concentration of a benzene solution prepared by mixing 12.0 g C6H6 with 38.0 g CCl4 is ______
molal.
A) 4.04
B) 0.240
C) 0.622
D) 0.316
E) 0.508
149. Which produces the greatest number of ions when one mole dissolves in water?
A) NaCl
B) NH4NO3
C) NH4Cl
D) Na2SO4
E) C12H22O11
150. Of the following, a 0.1 M aqueous solution of __________ will have the lowest freezing point.
A) NaCl
B) Al(NO3)3
C) K2CrO4
D) Na2SO4
E) sucrose, C12H22O11
151. Colligative properties of solutions include all of the following except __________.
A) depression of vapor pressure upon addition of a solute to a solvent
B) elevation of the boiling point of a solution upon addition of a solute to a solvent
C) depression of the freezing point of a solution upon addition of a solute to a solvent
D) an increase in the osmotic pressure of a solution upon the addition of more solute
E) the increase of reaction rates with increase in temperature
Acids and Bases
152. What is the conjugate acid of NH3?
A) NH3
D) NH4+
B) NH2+
E) NH4OH
153. The conjugate base of HSO4- is __________.
A) OHB) H2SO4
D) HSO4+
E) H3SO4+
C) NH3+
C) SO4 2-
154. According to the reaction H2O + H2SO4 → H3O+ + HSO4-, which molecule is acting as base?
A) H2SO4
D) HSO4-
B) H2O
E) None of the above
C) H3O+
155. According to the reaction H3O+ + HSO4- → H2O + H2SO4, which molecule is acting as an acid?
A) H2SO4
B) H2O
C) H3O+
D) HSO4
E) None of the above
156. A substance that is capable of acting as both an acid and as a base is __________.
A) autosomal
B) conjugated
C) amphoteric
D) saturated
E) miscible
157. The molar concentration of hydronium ion in pure water at 25 °C is __________.
A) 0.00
B) 1.0  10 7
C) 1.0 1014
D) 1.00
E) 7.00
158. Nitric acid is a strong acid. This means that __________.
A) aqueous solutions of HNO3 contain equal concentrations of H+(aq) and OH-(aq)
B) HNO3 does not dissociate at all when it is dissolved in water
C) HNO3 dissociates completely to H+(aq) and NO3-(aq) when it dissolves in water
D) HNO3 produces a gaseous product when it is neutralized
E) HNO3 cannot be neutralized by a weak base
159. If you had a 1.0 M solution of a strong acid, what would be a reasonable pH?
A) 1 B) 6
C) 7
D) 8 E) 13
160. If you had a 1.0 M solution of a weak base, what would be a reasonable pH?
A) 1 B) 6
C) 7
D) 8 E) 13
161. As the pH increases, the hydroxide ion concentration __________.
A) decreases
B) increases
C) starts to affect the [H+]
D) stays constant
162. If the pH of a solution was 7 and you were to increase the hydronium ion concentration 1000x, what
would the pH be?
A) 4 B) 7
C) 143
D) 7000
E) 1 x 10-4
163. What is the pH of an aqueous solution at 25.0 °C that contains 3.98 x 10-9 M hydronium ion?
A) 8.40
B) 5.60
C) 9.00
D) 3.98
E) 7.00
164. What is the concentration (in M) of hydronium ions in a solution at 25.0 °C with pH = 4.282?
A) 4.28
B) 9.71
C) 1.92 x 10-10
D) 5.22 x 10-5
E) 1.66 x 104
Kinetics and Equilibrium
165. Which of the following will not affect the rate of a chemical reaction?
A) increasing the concentration of a reactant
B) decreasing the concentration of a reactant
C) increasing the temperature at which the reaction is carried out
D) performing the reaction in the presence of a catalyst
E) All of the above will affect the rate of a chemical reaction.
166. A reaction was found to be second order in carbon monoxide concentration. The rate of the reaction
__________ if the [CO] is doubled, with everything else kept the same.
A) doubles
B) remains unchanged
C) triples
D) increases by a factor of 4
E) is reduced by a factor of 2
167. If the rate law for a reaction is first order in A and second order in B, then the rate law is
A) k[A][B]
B) k[A]2[B]3
C) k[A][B]2
D) k[A]2[B]
E) k[A]2[B]2
168. The rate law for a reaction is rate = k[A][B]2. Which one of the these statements is false?
A) The reaction is first order in A.
B) The reaction is second order in B.
C) The reaction is second order overall.
D) k is the reaction rate constant
E) If [B] is doubled, the reaction rate will increase by a factor of 4.
169. Which energy difference in this energy profile corresponds to the activation energy for the forward
reaction?
A) x
B) y
C) x+y
D) x-y
E) y-x
170. In the diagram, what quantity is represented by “ y” ?
A) heat of reaction, ΔH
B) activation energy of the forward reaction
C) activation energy of the reverse reaction
D) potential energy of the reactants
E) potential energy of the products
171. In the energy profile of a reaction, the species that exists at
the maximum on the curve is called the
A) product.
B) activated
complex.
C) activation
energy.
D) enthalpy of reaction.
E) atomic state.
172.
A)
B)
C)
D)
A catalyst can increase the rate of a reaction __________.
by lowering the activation energy of the reverse reaction
by increasing the overall activation energy (Ea) of the reaction
by providing an alternative pathway with a lower activation energy
All of these are ways that a catalyst might act to increase the rate of reaction.
173. Which of the following expressions is the correct equilibrium-constant expression for the reaction below?
HF(aq) + H2O (l) → H3O+(aq) + F-(aq)
A) [HF] [H2O] / [H3O+] [F-]
B) 1 / [HF]
C) [F-] / [HF]
+
+
D) [H3O ] [F ] / [HF]
E) [H3O ] [F ] / [HF] [H2O]
174.
A)
B)
C)
D)
E)
The equilibrium constant for a reaction is Keq = 4.34 x 10-3 at 300 °C. At equilibrium, __________.
products predominate
reactants predominate
roughly equal amounts of products and reactants are present
only products are present
only reactants are present
175. Of the following equilibria, only __________ will shift to the left in response to a decrease in volume.
A) H2(g) + Cl2(g) → 2 HCl(g)
B) 2 SO3(g) → 2 SO2(g) + O2(g)
C) N2(g) + 3 H2(g) → 2 NH3(g)
D) 4 Fe(s) + 3 O2(g) → 2 Fe2O3(s)
E) 2 HI(g) → H2(g) + 2 I2(g)
176. The reaction 2 SO2(g) + O2(g) → 2 SO3(g) is exothermic. Le Chatelier's Principle predicts that __________
will result in an increase in the number of moles of SO 3(g) in the reaction container.
A) increasing the pressure
B) decreasing the pressure
C) increasing the temperature
D) removing some oxygen
E) increasing the volume of the container
177. Consider the reaction 2 NH3(g) → 2 N2(g) + 3 H2(g) (ΔH° = +92.4 kJ) at equilibrium. Le Chatelier's principle
predicts that adding N2(g) to the system at equilibrium will result in __________.
A) a decrease in the concentration of NH3(g)
B) a decrease in the concentration of H2(g)
C) an increase in the value of the equilibrium constant
D) a lower partial pressure of N2(g)
E) removal of all of the H2(g)
178.
A)
B)
C)
D)
E)
The effect of a catalyst on an equilibrium is to __________.
increase the rate of the forward reaction only
increase the equilibrium constant so that products are favored
slow the reverse reaction only
shift the equilibrium to the right
increase the rate at which equilibrium is achieved without changing the composition of the equilibrium
mixture