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Chemistry 400: General Chemistry
Discussion: Review for Multiple Choice Section of Final
Chapter 3
1) Which of the following contains BOTH ionic and covalent
bonds?
A) CaI2
B) COS
C) CaSO4
D) SF6
E) None of the above contain both ionic and covalent
bonds.
5) What is the mass (in kg) of 6.89 × 1025 molecules of CO2?
The molar mass of CO2 is 44.01 g/mol.
A) 3.85 kg
B) 5.04 kg
C) 2.60 kg
D) 3.03 kg
E) 6.39 kg
6) Calculate the mass percent composition of lithium in Li3PO4.
A) 26.75 %
B) 17.98 %
C) 30.72 %
D) 55.27 %
E) 20.82 %
2) Write the formula for barium nitrite.
A) Ba3N2
B) BaNO3
C) BN
D) Ba(NO2)2
E) B(NO2)3
7) Determine the molecular formula of a compound that is
49.48% carbon, 5.19% hydrogen, 28.85% nitrogen, and 16.48%
oxygen. The molecular weight is 194.19 g/mol.
A) C8H12N4O2
B) C4H5N2O
C) C8H10N4O2
D) C8H10N2O
3) What is the charge on the Cr ions in Cr2O3?
A) 2B) 1+
C) 2+
D) 3+
4) Calculate the molar mass of Al(C2H3O2)3.
A) 86.03 g/mol
B) 204.13 g/mol
C) 56.00 g/mol
D) 258.09 g/mol
E) 139.99 g/mol
8) Write a balanced equation to show the reaction of sulfurous
acid with lithium hydroxide to form water and lithium sulfite.
A) H2SO4 (aq) + LiOH (aq) → H2O (l) + Li2SO4 (aq)
B) H2SO3 (aq) + 2 LiOH (aq) → 2 H2O (l) + Li2SO3
(aq)
C) HSO3 (aq) + LiOH (aq) → H2O (l) + LiSO3 (aq)
D) HSO4 (aq) + LiOH (aq) → H2O (l) + LiSO4 (aq)
E) H2S (aq) + 2 LiOH (aq) → 2 H2O (l) + Li2S (aq)
Chapter 4
2) How many molecules of H2S are required to form 79.0 g of
sulfur according to the following reaction? Assume excess SO2.
1) According to the following balanced reaction, how many
moles of HNO3 are formed from 8.44 moles of NO2 if there is
plenty of water present?
2 H2S(g) + SO2(g) → 3 S(s) + 2H2O(l)
A) 1.48 × 1024 molecules H2S
B) 9.89 × 1023 molecules H2S
3 NO2(g) + H2O(l) → 2 HNO3(aq) + NO(g)
C) 5.06 × 1025 molecules H2S
D) 3.17 × 1025 molecules H2S
A) 2.81 moles HNO3
B) 25.3 moles HNO3
C) 8.44 moles HNO3
D) 5.63 moles HNO3
E) 1.83 moles HNO3
E) 2.44 × 1023 molecules H2S
96
3) Determine the theoretical yield of H2S (in moles) if 4.0 mol
Al2S3 and 4.0 mol H2O are reacted according to the following
balanced reaction.
9) What volume of 0.305 M AgNO3 is required to react exactly
with 155.0 mL of 0.274 M Na2SO4 solution? Hint: You will
want to write a balanced reaction.
A) 581 mL
B) 173 mL
C) 345 mL
D) 139 mL
E) 278 mL
Al2S3 (s) + 6 H2O(l) → 2 Al(OH)3 (s) + 3 H2S(g)
A) 12 mol H2S
B) 4.0 mol H2S
C) 18 mol H2S
D) 6.0 mol H2S
E) 2.0 mol H2S
10) Which of the following solutions will have the highest
electrical conductivity?
A) 0.045 M Al2(SO4)3
B) 0.050 M (NH4)2CO3
C) 0.10 M LiBr
D) 0.10 M NaI
E) 0.10 M KF
4) How many moles of LiI are contained in 258.6 mL of 0.0296
M LiI solution?
A) 1.31 × 10-3 mol
B) 8.74 × 10-3 mol
C) 1.14 × 10-3 mol
D) 3.67 × 10-3 mol
E) 7.65 × 10-3 mol
11) Identify the spectator ions in the following molecular
equation.
KBr(aq) + AgNO3(aq) → AgBr(s) + KNO3(aq)
5) What is the concentration of magnesium ions in a 0.125 M
Mg(NO3)2 solution?
A) 0.125 M
B) 0.0625 M
C) 0.375 M
D) 0.250 M
E) 0.160 M
A) Ag+ and BrB) K+ and NO3C) K+ and BrD) Ag+ and NO3E) There are no spectator ions in this reaction.
6) What volume (in mL) of 0.0887 M MgF2 solution is needed to
make 275.0 mL of 0.0224 M MgF2 solution?
A) 72.3 mL
B) 91.8 mL
C) 10.9 mL
D) 69.4 mL
E) 14.4 mL
12) Give the net ionic equation for the reaction (if any) that
occurs when aqueous solutions of Na2CO3 and HCl are mixed.
A) 2 H+(aq) + CO32-(aq) → H2CO3(s)
B) 2 Na+(aq) + CO32-(aq) + 2 H+(aq) + 2 Cl-(aq) →
H2CO3(s) + 2 NaCl(aq)
C) 2 H+(aq) + CO32-(aq) → H2O(l) + CO2(g)
D) 2 Na+(aq) + CO32-(aq) + 2 H+(aq) + 2 Cl-(aq) →
H2CO3(s) + 2 Na+(aq) + 2 Cl-(aq)
7) Determine the molarity of a solution formed by dissolving 468
mg of MgI2 in enough water to yield 50.0 mL of solution.
A) 0.0297 M
B) 0.0337 M
C) 0.0936 M
D) 0.0107 M
E) 0.0651 M
E) No reaction occurs.
13) Determine the oxidation state of S in MgSO4.
A) -4
B) +2
C) +4
D) +6
E) -2
8) How many ions are present in 30.0 mL of 0.600 M Na2CO3
solution?
A) 3.25 × 1022 ions
B) 2.17 × 1022 ions
C) 1.08 × 1022 ions
D) 5.42 × 1022 ions
E) 3.61 × 1021 ions
14) What element is undergoing oxidation (if any) in the
following reaction?
CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g)
A) O
B) H
C) C
D) both C and H
E) None of the elements is undergoing oxidation.
97
Chapter 5
8) Using the graph below, determine the gas that has the lowest
density at STP.
1) The atmospheric pressure is 715 mm Hg. What is the pressure
in inches of Hg?
A) 23.9 in Hg
B) 0.940 in Hg
C) 48.6 in Hg
D) 28.1 in Hg
E) 30.0 in Hg
2) A gas occupies 3.33 L at 2.23 atm. What is the volume at 2.50
atm?
A) 1.67 L
B) 3.73 L
C) 2.97 L
D) 0.268 L
E) 18.6 L
3) A sample of 0.300 moles of nitrogen occupies 0.600 L. Under
the same conditions, what number of moles occupies 1.200 L?
A) 0.600 moles
B) 1.50 moles
C) 0.33 moles
D) 6.00 moles
A) A
B) B
C) C
D) D
E) All of the gases have the same density at STP.
9) Define vapor pressure.
A) partial pressure of water in a liquid mixture
B) partial pressure of water in a gaseous mixture
C) condensation of water
D) water dissolved in a liquid
E) water molecules
4) A syringe initially holds a sample of gas with a volume of 285
mL at 355 K and 1.88 atm. To what temperature must the gas in
the syringe be heated/cooled in order to have a volume of 435 mL
at 2.50 atm?
A) 139 K
B) 572 K
C) 175 K
D) 466 K
E) 721 K
10) A mixture of 0.220 moles CO, 0.350 moles H2 and 0.640
moles He has a total pressure of 2.95 atm. What is the pressure
of CO?
A) 1.86 atm
B) 0.649 atm
C) 0.536 atm
D) 1.54 atm
E) 0.955 atm
5) What is the volume of 9.783 x 1023 atoms of He at 9.25 atm
and 512K?
A) 7.38 L
B) 3.69 L
C) 1.85 L
D) 15.4 L
E) 30.8 L
11) The following reaction is used to generate hydrogen gas in
the laboratory. If 243 mL of gas is collected at 25 °C and has a
total pressure of 745 mm Hg, what mass of hydrogen is
produced? A possibly useful table of water vapor pressures is
provided below.
6) The density of a gas is 1.43 g/L at 1.00 atm and 273 K. What
is the gas?
A) Cl2
B) S
C) O2
D) Ne
Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g)
T (°C)
20
25
30
7) A compound is found to be 30.45% N and 69.55 % O by mass.
If 1.63 g of this compound occupy 389 mL at 0.00°C and 775
mm Hg, what is the molecular formula of the compound?
A) NO2
B) N2O
C) N4O2
D) N2O5
E) N2O4
P (mm Hg)
17.55
23.78
31.86
A) 0.0196 g H2
B) 0.0717 g H2
C) 0.0190 g H2
D) 0.0144 g H2
E) 0.0449 g H2
98
13) Which of the following statements is TRUE?
A) At a given temperature, lighter gas particles travel
more slowly than heavier gas particles.
B) The smaller a gas particle, the slower it will effuse
C) The higher the temperature, the lower the average
kinetic energy of the sample.
D) At low temperatures, intermolecular forces become
important and the pressure of a gas will be lower than
predicted by the ideal gas law.
E) None of the above statements are true.
12) Determine the volume of O2 (at 1.00 atm and 273 K) formed
when 50.0 g of KClO3 decomposes according to the following
reaction.
2 KClO3(s) → 2 KCl(s) + 3 O2(g)
A) 9.14 L
B) 8.22 L
C) 12.3 L
D) 13.7 L
E) 14.6 L
Chapter 6
5) According to the following reaction, how much energy is
required to decompose 55.0 kg of Fe3O4?
1) Which of the following (with specific heat capacity provided)
would show the smallest temperature change upon gaining 200.0
J of heat?
A) 50.0 g Al, CAl = 0.903 J/g°C
B) 50.0 g Cu, CCu = 0.385 J/g°C
C) 25.0 g granite, Cgranite = 0.79 J/g°C
D) 25.0 g Au, CAu = 0.128 J/g°C
E) 25.0 g Ag, CAg = 0.235 J/g°C
Fe3O4(s) → 3 Fe(s) + 2 O2(g)
ΔH°rxn = +1118 kJ
A) 1.10 × 106 kJ
B) 2.38 × 102 kJ
C) 2.66 × 105 kJ
D) 1.12 × 103 kJ
E) 3.44 × 104 kJ
2) Calculate the amount of heat (in kJ) required to raise the
temperature of a 79.0 g sample of ethanol from 298.0 K to 385.0
K. The specific heat capacity of ethanol is 2.42 J/g°C.
A) 57.0 kJ
B) 16.6 kJ
C) 73.6 kJ
D) 28.4 kJ
E) 12.9 kJ
6) Using the following thermochemical equation, determine the
amount of heat produced from the combustion of 24.3 g benzene
(C6H6).
2 C6H6(l) + 15 O2(g) → 12 CO2(g) + 6 H2O(g)
ΔH°rxn = -6278 kJ/mol
A) -3910 kJ
B) -1950 kJ
C) -977 kJ
D) -40.1 kJ
E) -0.302 kJ
3) Which of the following processes is exothermic?
A) the formation of dew in the morning
B) the melting of ice
C) the chemical reaction in a "cold pack" often used to
treat injuries
D) the vaporization of water
E) None of the above are exothermic.
7) According to the following reaction, how much energy is
evolved during the reaction of 32.5 g B2H6 and 72.5 g Cl2?
B2H6(g) + 6 Cl2(g) → 2 BCl3(g) + 6 HCl(g)
ΔH°rxn = -1396 kJ/mol
4) Using the following equation for the combustion of octane,
calculate the heat of reaction for 100.0 g of octane.
A) -1640 kJ
B) -238 kJ
C) -1430 kJ
D) -3070 kJ
E) -429 kJ
2 C8H18 + 25 O2 → 16 CO2 + 18 H2O
ΔH°rxn = -11018 kJ
A) -4.82 x 103 kJ/mol
B) -4.82 kJ/mol
C) -9.64 x 103 kJ/mol
D) -1.26 x 104 kJ/mol
99
ΔH
8) According to the following reaction, how much energy is
evolved during the reaction of 2.50 L B2H6 and 5.65 L Cl2 (Both
gases are initially at STP)?
11) Use the ΔH°f and ΔH°rxn information provided to calculate
ΔH°f for IF:
IF7(g) + I2(g) → IF5(g) + 2 IF(g)
ΔH°rxn = -1396 kJ
B2H6(g) + 6 Cl2(g) → 2 BCl3(g) + 6 HCl(g)
ΔH°rxn = -1396 kJ/mol
A) -58.7 kJ
B) -156 kJ
C) -215 kJ
D) -352 kJ
E) -508 kJ
IF7(g)
IF5(g)
9) Two aqueous solutions are both at room temperature and are
then mixed in a coffee cup calorimeter. The reaction causes the
temperature of the resulting solution to fall below room
temperature. Which of the following statements is TRUE?
A) The products have a lower potential energy than the
reactants.
B) This type of experiment will provide data to calculate
ΔErxn.
C) The reaction is exothermic.
D) Energy is leaving the system during reaction.
E) None of the above statements are true.
ΔH°f
(kJ/mol)
-941
-840
A) 101 kJ/mol
B) -146 kJ/mol
C) -190. kJ/mol
D) -95 kJ/mol
E) 24 kJ/mol
12) Use the information provided to determine ΔH°rxn for the
following reaction:
9) A 100.0 mL sample of 0.300 M NaOH is mixed with a 100.0
mL sample of 0.300 M HNO3 in a coffee cup calorimeter. If
both solutions were initially at 35.00°C and the temperature of
the resulting solution was recorded as 37.00°C, determine the
ΔH°rxn (in units of kJ/mol NaOH) for the neutralization reaction
between aqueous NaOH and HCl. Assume 1) that no heat is lost
to the calorimeter or the surroundings, and 2) that the density and
the heat capacity of the resulting solution are the same as water.
A) -55.7 kJ/mol NaOH
B) -169 kJ/mol NaOH
C) -16.7 kJ/mol NaOH
D) -27.9 kJ/mol NaOH
E) - 34.4 kJ/mol NaOH
CH4(g) + 3 Cl2(g) → CHCl3(l) + 3 HCl(g) ΔH°rxn = ?
CH4(g)
CHCl3(l)
HCl(g)
A) -151 kJ
B) -335 kJ
C) +662 kJ
D) +117 kJ
E) -217 kJ
10) Use the standard reaction enthalpies given below to
determine ΔH°rxn for the following reaction:
P4(g) + 10 Cl2(g) → 4PCl5(s)
ΔH°rxn = ?
Given:
PCl5(s) → PCl3(g) + Cl2(g)
ΔH°rxn = +157 kJ/mol
P4(g) + 6 Cl2(g) → 4 PCl3(g)
ΔH°rxn = -1207 kJ/mol
ΔH°rxn = -89 kJ
A) -1835 kJ
B) -1364 kJ
C) -1050. kJ
D) -1786 kJ
E) -2100. kJ
100
ΔH°f
(kJ/mol)
-75
-134
-92
Chapter 7
1) The vertical height of a wave is called
A) wavelength
B) amplitude
C) frequency
D) area
E) median
2) Which of the following visible colors of light has the longest
wavelength?
A) blue
B) green
C) yellow
D) red
E) violet
3) Which of the following occur as the energy of a photon
increases?
A) the frequency decreases.
B) the speed increases.
C) the wavelength increases
D) the wavelength gets shorter.
E) None of the above occur as the energy of a photon
increases.
4) Calculate the energy of the orange light emitted, per photon,
by a neon sign with a frequency of 4.89 × 1014 Hz.
A) 3.09 × 10-19 J
B) 6.14 × 10-19 J
C) 3.24 × 10-19 J
D) 1.63 × 10-19 J
E) 5.11 × 10-19 J
5) How many photons are contained in a flash of green light (525
nm) that contains 189 kJ of energy?
A) 5.67 × 1023 photons
B) 2.01 × 1024 photons
C) 1.25 × 1031 photons
D) 4.99 × 1023 photons
E) 7.99 × 1030 photons
6) Calculate the wavelength of a baseball (m = 155 g) moving at
32.5 m/s.
A) 7.60 × 10-36 m
B) 1.32 × 10-34 m
C) 2.15 × 10-32 m
D) 2.68 × 10-34 m
E) 3.57 × 10-32 m
7) Determine the velocity of a marble (m = 8.66 g) with a
wavelength of 3.46 × 10-33 m.
A) 45.2 m/s
B) 11.3 m/s
C) 22.1 m/s
D) 38.8 m/s
E) 52.9 m/s
8) Choose the transition (in a hydrogen atom) below that
represents the absorption of the shortest wavelength photon.
A) n = 1 to n = 2
B) n = 2 to n = 3
C) n = 4 to n = 5
D) n = 6 to n = 3
E) n = 3 to n = 1
9) Which of the following statements is TRUE?
A) We can sometimes know the exact location and speed
of an electron at the same time.
B) All orbitals in a given atom are roughly the same
size.
C) Since electrons have mass, we must always consider
them to have particle properties and never wavelike
properties.
D) Atoms are roughly spherical because when all of the
different shaped orbitals are overlapped, they take on a
spherical shape.
E) All of the above are true.
10) Identify the correct values for a 2p sublevel.
A) n = 3, l = 1, ml = 0
B) n = 2, l = 1, ml = -2
C) n = 1, l = 0, ml = 0
D) n = 2, l = 1, ml = 0
E) n = 4, l = -1, ml = -2
11) Determine the energy change associated with the transition
from n=2 to n=5 in the hydrogen atom.
A) -2.18 × 10-19 J
B) +6.54 × 10-19 J
C) +4.58 × 10-19 J
D) -1.53 × 10-19 J
E) +3.76 × 10-19 J
12) Determine the end (final) value of n in a hydrogen atom
transition, if the electron starts in
n = 4 and the atom emits a photon of light with a wavelength of
486 nm.
A) 1
B) 5
C) 3
D) 4
E) 2
13) What is the maximum number of f orbitals that are possible?
A) 1
B) 3
C) 7
D) 5
E) 9
101
Chapter 8
6) Place the following elements in order of decreasing atomic
radius.
1) Choose the orbital diagram that represents the ground state of
N.
A)
Xe
D)
Ar
A) Ar > Xe > Rb
B) Xe > Rb > Ar
C) Ar > Rb > Xe
D) Rb > Xe > Ar
E) Rb > Ar > Xe
B)
C)
Rb
7) Of the following, which atom has the smallest atomic radius?
A) K
B) As
C) Rb
D) Sb
8) Place the following in order of increasing radius.
Ca2+
E)
2) The element that corresponds to the electron configuration
1s22s22p6 is __________.
A) sodium
B) magnesium
C) lithium
D) beryllium
E) neon
3) The complete electron configuration of argon, element 18, is
__________.
A) 1s22s22p63s23p6
B) 1s22s22p103s23p2
C) 1s42s42p63s4
D) 1s42s42p10
E) 1s62s 62p23s4
4) Give the ground state electron configuration for Se.
A) [Ar]4s23d104p4
B) [Ar]4s24d104p4
C) [Ar]4s23d104p6
D) [Ar]4s23d10
E) [Ar]3d104p4
S2-
Cl-
A) Ca2+ < Cl- < S2B) Cl- < Ca2+ < S2C) S2- < Cl- < Ca2+
D) Ca2+ < S2- < ClE) Cl- < S2- < Ca2+
9) Which reaction below represents the second ionization of Sr?
A) Sr(g) → Sr+ (g) + eB) Sr2+ (g) + e- → Sr- (g)
C) Sr+ (g) + e- → Sr(g)
D) Sr- (g) + e- → Sr2- (g)
E) Sr+ (g) → Sr2+ (g) + e10) Choose the ground state electron configuration for Ti2+.
A) [Ar]3d2
B) [Ar]4s2
C) [Ar]4s23d2
D) [Ar]4s23d4
E) [Ar]3d4
5) How many unpaired electrons are present in the ground state
Kr atom?
A) 1
B) 2
C) 0
D) 3
E) 5
102
Chapter 9
7) Use the data given below to construct a Born-Haber cycle to
determine the lattice energy of CaO.
1) Which of the following represent the Lewis structure for N?
A)
Ca(s) → Ca(g)
Ca(g) → Ca+(g) + eCa+ (g) → Ca2+ (g) + e2 O(g) → O2(g)
O(g) + e- → O- (g)
O- (g) + e- → O2- (g)
B)
C)
D)
E)
2) Which of the following statements is TRUE?
A) An ionic bond is much stronger than most covalent
bonds.
B) An ionic bond is formed through the sharing of
electrons.
C) Ionic compounds at room temperature typically
conduct electricity.
D) Once dissolved in water, ionic compounds rarely
conduct electricity.
E) None of the above are true.
3) Identify the compound with the highest magnitude of lattice
energy.
A) NaCl
B) KCl
C) LiCl
D) CsCl
Ca(s) +
A) -3414 kJ
B) +1397 kJ
C) -2667 kJ
D) +3028 kJ
E) -2144 kJ
8) Choose the bond below that is most polar.
A) H-I
B) H-Br
C) H-F
D) H-Cl
E) C-H
9) Choose the best Lewis structure for CH2Cl2.
A)
4) Choose the compound below that should have the highest
melting point according to the ionic bonding model.
A) AlN
B) MgO
C) NaF
D) CaS
E) RbI
B)
C)
5) Give the complete electronic configuration for Ca2+.
A) 1s22s22p63s24p6
B) 1s22s22p63s23p6
C) 1s22s22p63s23p5
D) 1s22s23p64s25p6
E) 1s22s2p63s2p6
D)
E)
6) Place the following elements in order of decreasing
electronegativity.
S
Cl
1
O2(g) → CaO(s)
2
ΔH°(kJ/mol)
193
590
1010
-498
-141
878
-635
Se
A) Se > S > Cl
B) Cl > Se > S
C) Se > Cl > S
D) S > Cl > Se
E) Cl > S > Se
103
10) Choose the best Lewis structure for SF4.
A)
14) Choose the best Lewis structure for SO42 .
A)
B)
B)
C)
C)
D)
D)
E)
E)
11) Which of the following resonance structures for OCN- will
contribute most to the correct structure
of OCN-?
A) O(2 lone pairs)=C=N (2 lone pairs)
B) O(1 lone pair)≡C-N(3 lone pairs)
C) O(1 lone pair)=C(2 lone pairs)=N(1 lone pair)
D) O(3 lone pairs)-C≡N(with 1 lone pair)
E) They all contribute equally to the correct structure of
OCN-.
12) Which of the following elements can form compounds with
an expanded octet?
A) N
B) Br
C) F
D) Be
E) None of the above can form compounds with an
expanded octet.
13) Place the following in order of decreasing bond length.
H-F
H-I
H-Br
A) H-F > H-Br > H-I
B) H-I > H-F > H-Br
C) H-I > H-Br > H-F
D) H-Br > H-F > H-I
E) H-F > H-I > H-Br
15) Use the bond energies provided to estimate ΔH°rxn for the
reaction below.
CH3OH(l) + 2 O2(g) → CO2(g) + 2 H2O(g)
Bond
Bond Energy
(kJ/mol)
C-H
414
C-O
360
C=O
799
O=O
498
O-H
464
A) +473 kJ
B) -91 kJ
C) -486 kJ
D) -392 kJ
E) +206 kJ
16) Which compound has the highest carbon-carbon bond
strength?
A) CH3CH3
B) CH2CH2
C) HCCH
D) all bond strengths are the same
104
ΔH
Chapter 10
1) Give the approximate bond angle for a molecule with a
trigonal planar shape.
A) 109.5°
B) 180°
C) 120°
D) 105°
E) 90°
2) Give the approximate bond angle for a molecule with an
octahedral shape.
A) 109.5°
B) 180°
C) 120°
D) 105°
E) 90°
3) Determine the electron geometry (eg) of CO32-.
A) eg=linear
B) eg=tetrahedral
C) eg=trigonal planar
D) eg=trigonal bipyramidal
E) eg=octahedral
4) Determine the electron geometry (eg) of NCl3.
A) eg=linear
B) eg=tetrahedral
C) eg=trigonal planar
D) eg=trigonal bipyramidal
E) eg=octahedral
5) Determine the electron geometry (eg) of XeF4.
A) eg=linear
B) eg=tetrahedral
C) eg=trigonal planar
D) eg=trigonal bipyramidal
E) eg=octahedral
6) Place the following in order of increasing F-A-F bond angle,
where A represents the central atom in each molecule.
PF3
OF2
PF43-
A) PF3 < OF2 < PF43B) OF2 < PF3 < PF43C) OF2 < PF43- < PF3
D) PF43- < OF2 < PF3
E) PF43- < PF3 < OF2
7) A molecule containing a central atom with sp3d hybridization
has a(n) __________ electron geometry.
A) tetrahedral
B) linear
C) octahedral
D) trigonal planar
E) trigonal bipyramidal
8) Place the following in order of decreasing X-A-X bond angle,
where A represents the central atom and X represents the outer
atoms in each molecule.
N 2O
NCl3
NO2-
A) NCl3 > NO2- > N2O
B) NO2- > N2O > NCl3
C) N2O > NO2- > NCl3
D) NCl3 > N2O > NO2E) N2O > NCl3 > NO29) How many of the following molecules are polar?
XeCl2
COF2
PCl4F
SF6
A) 0
B) 3
C) 1
D) 2
E) 4
10) Determine the electron geometry and polarity of SF6 .
A) eg=trigonal bipyramidal, nonpolar
B) eg=tetrahedral, polar
C) eg=trigonal bipyramidal, polar
D) eg=octahedral, nonpolar
E) eg=octahedral, polar
11) Determine the electron geometry and polarity of HBrO2 .
A) eg=trigonal bipyramidal, nonpolar
B) eg=octahedral, nonpolar
C) eg=tetrahedral, polar
D) eg=tetrahedral, nonpolar
E) eg=linear, polar
12) Choose the compound below that contains at least one polar
covalent bond, but is nonpolar.
A) GeH2Br2
B) SCl2
C) AsCl5
D) CF2Cl2
E) All of the above are nonpolar and contain a polar
covalent bond.
13) Describe a pi bond.
A) side by side overlap of p orbitals
B) end to end overlap of p orbitals
C) s orbital overlapping with the end of a p orbital
D) overlap of two s orbitals
E) p orbital overlapping with a d orbital
105
14) Use the molecular orbital diagram shown to determine which
of the following are paramagnetic.
17) Use the molecular orbital diagram shown to determine which
of the following is most stable.
A) B22+
B) B22-
A) C22B) N22-
C) N22+
D) C22-
C) B2
D) C22E) B22-
E) B2
15) Determine the electron geometry (eg) of the underlined
carbon in CH3CN.
A) eg=tetrahedral
B) eg=linear
C) eg=trigonal planar
D) eg=trigonal bipyramidal
E) eg=octahedral
16) Draw the Lewis structure for H3O+. What is the
hybridization on the O atom?
A) sp
B) sp3
C) sp2
D) sp3d
E) sp3d2
18) Draw the molecular orbital diagram shown to determine
which of the following is paramagnetic.
A) B22+
B) B22C) N22D) C22E) B2
19) Draw the Lewis structure for BrF5. What is the hybridization
on the Br atom?
A) sp3d2
B) sp3d
C) sp3
D) sp2
E) sp
106
20) How many of the following molecules have sp3 hybridization
on the central atom?
XeCl4
CH4
SF4
21) How many of the following molecules have sp3d2
hybridization on the central atom?
C 2H 2
SeCl6
A) 0
B) 4
C) 3
D) 2
E) 1
Chapter 11
3) What type of intermolecular force causes the dissolution of
NaCl in water?
A) hydrogen bonding
B) dipole-dipole forces
C) ion-dipole force
D) dispersion forces
E) none of the above
4) Choose the molecule or compound that exhibits dipole-dipole
forces as its strongest intermolecular force.
A) H2
B) SO2
C) NH3
D) CF4
E) BCl3
IF5
AsCl5
A) 1
B) 3
C) 0
D) 2
E) 4
6) Place the following compounds in order of increasing strength
of intermolecular forces.
1) Which of the following statements is TRUE?
A) Intermolecular forces are generally stronger than
bonding forces.
B) The potential energy of molecules decrease as they
get closer to one another.
C) Energy is given off when the attraction between two
molecules is broken.
D) Increasing the pressure on a solid usually causes it to
become a liquid.
E) None of the above are true.
2) What is the strongest type of intermolecular force present in
CHF3?
A) ion-dipole
B) dispersion
C) hydrogen bonding
D) dipole-dipole
E) none of the above
XeF4
CH4
CH3CH2CH3
CH3CH3
A) CH3CH2CH3 < CH4 < CH3CH3
B) CH3CH2CH3 < CH3CH3 < CH4
C) CH3CH3 < CH4 < CH3CH2CH3
D) CH4 < CH3CH2CH3 < CH3CH3
E) CH4 < CH3CH3 < CH3CH2CH3
7) Choose the pair of substances that are most likely to form a
homogeneous solution.
A) C6H14 and C10H20
B) LiBr and C5H12
C) N2O4 and NH4Cl
D) C6H14 and H2O
E) None of the pairs above will form a homogeneous
solution.
8) Which of the following statements is FALSE?
A) The rate of vaporization increases with increasing
surface area.
B) The rate of vaporization increases with decreasing
strength of intermolecular forces.
C) The rate of vaporization increases with increasing
temperature.
D) Molecules with hydrogen bonding are more volatile
than compounds with dipole-dipole forces.
E) None of the above are false.
9) Place the following substances in order of increasing boiling
point.
CH3CH2OH
Ar
CH3OCH3
A) Ar < CH3OCH3 < CH3CH2OH
B) CH3CH2OH < Ar < CH3OCH3
C) CH3CH2OH < CH3OCH3 < Ar
D) CH3OCH3 < Ar < CH3CH2OH
E) Ar < CH3CH2OH < CH3OCH3
5) Identify the compound that does not have hydrogen bonding.
A) (CH3)3N
B) H2O
C) CH3OH
D) HF
E) CH3NH2
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10) Which of the following substances would you predict to have
the highest ΔHvap?
A) Xe
B) C6H6
C) SiF4
D) Br2
E) N2
11) How much energy is required to vaporize 158 g of butane
(C4H10) at its boiling point, if its ΔHvap is 24.3 kJ/mol?
A) 15.1 kJ
B) 66.1 kJ
C) 89.4 kJ
D) 11.2 kJ
E) 38.4 kJ
12) Determine the normal boiling point of a substance whose
vapor pressure is 55.1 mm Hg at 35°C and has a ΔHvap of 32 .1
kJ/mol.
A) 255 K
B) 368 K
C) 412 K
D) 390. K
E) 466 K
13) Ethanol (C2H5OH) melts at -114°C. The enthalpy of fusion
is 5.02 kJ/mol. The specific heats of solid and liquid ethanol are
0.97 J/gK and 2.3 J/gK, respectively. How much heat (kJ) is
needed to convert 25.0 g of solid ethanol at -135°C to liquid
ethanol at -50°C?
A) 207.3 kJ
B) -12.7 kJ
C) 6.91 kJ
D) 4192 kJ
E) 9.21 kJ
Chapter 12
1) Dynamic equilibrium can be defined as
A) rate of dissolution = rate of deposition
B) rate of dissolution < rate of deposition
C) rate of dissolution > rate of deposition
D) rate of bubbling > rate of dissolving
E) rate of evaporating > rate of condensing
2) What mass (in g) of NH3 must be dissolved in 475 g of
methanol to make a 0.250 m solution?
A) 2.02 g
B) 4.94 g
C) 1.19 g
D) 8.42 g
E) 1.90 g
14) Based on the figure above, the boiling point of water under an
external pressure of 0.316 atm is __________°C.
A) 70
B) 40
C) 60
D) 80
E) 90
3) Determine the molality of a solution prepared by dissolving
0.500 moles of CaF2 in 11.5 moles H2O.
A) 1.88 m
B) 4.35 m
C) 5.31 m
D) 4.14 m
E) 2.41 m
4) A solution is prepared by dissolving 98.6 g of NaCl in enough
water to form 875 mL of solution. Calculate the mass % of the
solution if the density of the solution is 1.06 g/mL.
A) 11.3%
B) 12.7%
C) 9.4%
D) 10.6%
E) 11.9%
108
5) Determine the freezing point of a solution that contains 78.8 g
of naphthalene (C10H8, molar mass = 128.16 g/mol) dissolved in
722 mL of benzene (d = 0.877 g/mL). Pure benzene has a
melting point of 5.50°C and a freezing point depression constant
of -4.90°C/m.
A) 4.76°C
B) 4.17°C
C) 0.74°C
D) 1.33°C
E) 1.68°C
8) A 150.0 mL sample of an aqueous solution at 25°C contains
15.2 mg of an unknown nonelectrolyte compound. If the solution
has an osmotic pressure of 8.44 torr, what is the molar mass of
the unknown compound?
A) 223 g/mol
B) 294 g/mol
C) 341 g/mol
D) 448 g/mol
E) 195 g/mol
9) Choose the aqueous solution that has the highest boiling point.
These are all solutions of nonvolatile solutes and you should
assume ideal van't Hoff factors where applicable.
A) 0.100 m NaNO3
B) 0.100 m Li2SO4
C) 0.200 m C3H8O3
D) 0.060 m Na3PO4
E) They all have the same boiling point.
6) Calculate the boiling point of a solution of 500.0 g of ethylene
glycol (C2H6O2) dissolved in 500.0 g of water. Kf = 1.86°C/m
and Kb = 0.512°C/m. Use 100°C as the boiling point of water.
A) 108°C
B) 92°C
C) 130°C
D) 70°C
E) 8.3°C
7) Determine the boiling point of a solution that contains 78.8 g
of naphthalene (C10H8, molar mass = 128.16 g/mol) dissolved in
722 mL of benzene (d = 0.877 g/mL). Pure benzene has a boiling
point of 80.1°C and a boiling point elevation constant of
2.53°C/m.
A) 2.2°C
B) 2.5°C
C) 82.3°C
D) 80.4°C
E) 82.6°C
10) Choose the aqueous solution below with the lowest freezing
point. These are all solutions of nonvolatile solutes and you
should assume ideal van't Hoff factors where applicable.
A) 0.075 m NaI
B) 0.075 m (NH4)3PO4
C) 0.075 m NaBrO4
D) 0.075 m LiCN
E) 0.075 m KNO2
109