Download ____ 1. The energy required to convert a ground

____
1. The energy required to convert a ground-state atom in the gas phase to a gaseous positive ion
a. Activation energy
d. Kinetic energy
b. Free energy
e. Lattice energy
c. Ionization energy
____
2. The energy change that occurs in the conversion of an ionic solid to widely separated gaseous ions
a. Activation energy
d. Kinetic energy
b. Free energy
e. Lattice energy
c. Ionization energy
____
3. Represents an atom that is chemically unreactive
a. 1s ___ 2s 
d.
b. 1s  2s 
e.
c. 1s  2s  2p   _
1s  2s  2p   
[Ar] 4s  3d     
4. Represents an atom in an excited state
a. 1s ___ 2s 
b. 1s  2s 
c. 1s  2s  2p   _
1s  2s  2p   
[Ar] 4s  3d     
____
d.
e.
____
5. Represents an atom that has four valence electrons
a. 1s ___ 2s 
d. 1s  2s  2p   
b. 1s  2s 
e. [Ar] 4s  3d     
c. 1s  2s  2p   __
____
6. Represents an atom of a transition metal
a. 1s ___ 2s 
b. 1s  2s 
c. 1s  2s  2p   __
____
____
____
____
7. Cesium chloride, CsCl(s)
a. Lattice of positive and negative ions held
together by electrostatic forces
b. Closely packed lattice with delocalized
electrons throughout
c. Strong single covalent bonds with weak
intermolecular forces
8. Gold, Au(s)
a. Lattice of positive and negative ions held
together by electrostatic forces
b. Closely packed lattice with delocalized
electrons throughout
c. Strong single covalent bonds with weak
intermolecular forces
9. Carbon dioxide, CO2(s)
a. Lattice of positive and negative ions held
together by electrostatic forces
b. Closely packed lattice with delocalized
electrons throughout
c. Strong single covalent bonds with weak
intermolecular forces
10. Methane, CH4(s)
a. Lattice of positive and negative ions held
together by electrostatic forces
b. Closely packed lattice with delocalized
electrons throughout
d.
e.
1s  2s  2p   
[Ar] 4s  3d     
d.
Strong multiple covalent bonds (including
p-bonds) with weak intermolecular forces
Macromolecules held together with strong
polar bonds
e.
d.
e.
d.
e.
d.
e.
Strong multiple covalent bonds (including
p-bonds) with weak intermolecular forces
Macromolecules held together with strong
polar bonds
Strong multiple covalent bonds (including
p-bonds) with weak intermolecular forces
Macromolecules held together with strong
polar bonds
Strong multiple covalent bonds (including
p-bonds) with weak intermolecular forces
Macromolecules held together with strong
polar bonds
Strong single covalent bonds with weak
intermolecular forces
c.
____
____
11. Is a gas in its standard state at 298 K
a. Lithium
b. Nickel
c. Bromine
d.
e.
Uranium
Flourine
12. Reacts with water to form a strong base
a. Lithium
b. Nickel
c. Bromine
d.
e.
Uranium
Flourine
____
13. What mass of Au is produced when 0.0500 mol of Au2S3 is reduced completely with excess H2?
a. 9.85 g
d. 39.4 g
b. 19.7 g
e. 48.9 g
c. 24.5 g
____
14. When a solution of sodium chloride is vaporized in a flame, the color of the flame is
a. blue
d. violet
b. yellow
e. white
c. green
____
15. A hot-air balloon rises. Which of the following is the best explanation for this observation?
a. The pressure on the walls of the balloon
d. The rate of diffusion of cooler air is less
increases with increasing temperature.
than that of warmer air.
b. The difference in temperature between the e. The air density inside the balloon is less
air inside and outside the balloon
than that of the surrounding air.
produces convection currents.
c. The cooler air outside the balloon pushes
in on the walls of the balloon.
____
16. The safest and most effective emergency procedure to treat an acid splash on skin is to do which of the
following immediately?
a. Dry the affected area with paper towels
d. Flush the affected area with water and
then with a dilute NaHCO3 solution
b. Sprinkle the affected area with powdered e. Flush the affected area with water and
Na2SO4(s)
then with a dilute vinegar solution
c. Flush the affected area with water and
then with a dilute NaOH solution
Temperature
P
R
S
Q
T
Time
____
17. The cooling curve for a pure substance as it changes from a liquid to a solid is shown above. The solid
and liquid coexist at
a. point Q only
d. all points on the curve between R and T
b. point R only
e. no point on the curve
c. all points on the curve between Q and S
____
18. ___C10H12O4S (s) + ___O2 (g)  ___CO2 (g) + ___SO2 (g) + ___H2O (g)
When the equation above is balanced and all coefficients are reduced to their lowest whole-number
terms, the coefficient for O2 (g) is
a. 6
d. 14
b. 7
e. 28
c. 12
____
19. Appropriate uses of a visible-light spectrophotometer include which of the following?
I. Determining the concentration of a solution of Cu(NO3)2
II. Measuring the conductivity of a solution of KMnO4
III. Determining which ions are present in a solution that may contain Na+, Mg2+, Al3+
a. I only
d. I and II only
b. II only
e. I and III only
c. III only
____
20. The melting point of MgO is higher than that of NaF. Explanations for this observation include which of
the following?
I. Mg2+ is more positively charged than Na+.
II. O2- is more negatively charged than F-.
III. The O2- ion is smaller than the F- ion.
a. II only
d. II and III only
b. I and II only
e. I, II, and III
c. I and III only
CH 3
O
||
 C  CH 2  CH 3
____
21. The organic compound represented above is an example of
a. an organic acid
d. an aldehyde
b. an alcohol
e. a ketone
c. an ether
____
22. H2Se (g) + 4 O2F2 (g)  SeF6 (g) + 2 HF (g) + 4 O2 (g)
Which of the following is true regarding the reaction represented above?
a. The oxidation number of O does not
d. The oxidation number of Se changes from
change.
-2 to +6.
b. The oxidation number of H changes from - e. It is a disproportionation reaction for F.
1 to +1.
c. The oxidation number of F changes from
+1 to -1.
____
23. If the temperature of an aqueous solution of NaCl is increased from 20oC to 90oC, which of the following
statements is true?
a. The density of the solution remains
d. The mole fraction of the solute decreases.
unchanged.
b. The molarity of the solution remains
e. The mole fraction of solute increases.
unchanged.
c. The molality of the solution remains
unchanged.
____
24. Types of hybridization exhibited by the C atoms in propene, CH3CHCH2, include which of the following?
I. sp
II. sp2
III. sp3
a. I only
d. II and III only
b. III only
e. I, II, and III
c. I and II only
____
25. A 1.0 L sample of an aqueous solution contains 0.10 mol of NaCl and 0.10 mol of CaCl2. What is the
minimum number of moles of AgNO3 that must be added to the solution in order to precipitate all of the Clas AgCl (s) ? (Assume that AgCl is insoluble.)
a.
b.
c.
0.10 mol
0.20 mol
0.30 mol
d.
e.
Initial [NO] (mol L-1)
0.40 mol
0.60 mol
Initial [O2] (mol L-1)
Initial Rate of
Formation of NO2
(mol L-1s-1)
1
0.10
0.10
2.5 X 10-4
2
0.20
0.10
5.0 X 10-4
3
0.20
0.40
8.0 X 10-3
The initial-rate data in the table above were obtained for the reaction represented below. What is the
experimental rate law for the reaction?
2 NO (g) + O2 (g)  NO2 (g)
a. Rate = k[NO][O2]
d. Rate = k[NO]2[O2]2
b. Rate = k[NO][O2]2
e.
Rate = k
____
c.
26.
Experiment
Rate = k[NO]2[O2]
____
27. Ionization Energies for the element X (kJ mol-1)
First
Second
Third
Fourth
Fifth
580
1,815
2,740
11,600
14,800
The ionization energies for the element X are listed in the the table above. On the basis of the data,
element X is most likely to be
a. Na
d. Si
b. Mg
e. P
c. Al
____
28. The phase diagram for a pure substance is shown above. Which point on the diagram corresponds to the
equilibrium between the solid and liquid phases at the normal melting point?
a. A
d. D
b. B
e. E
c. C
____
29. Of the following molecules, which has the largest dipole moment?
a. CO
d. HF
b. CO2
e. F2
c. O2
____
30. ___Li3N (s) + ___H2O (l)  ___Li+ (aq) + ___OH- (aq) + ___NH3 (g)
When the equation above is balanced and all coefficients reduced to lowest whole number terms, the
coefficient for OH- (aq) is
a. 1
d. 4
b. 2
e. 6
c. 3
____
31. A sample of 61.8 g of H3BO3, a weak acid, is dissolved in the 1,000 g of water to make a 1.0-molal
solution. Which of the following would be the best procedure to determine the molarity of the solution?
(Assume no additional information is available.)
a. Titration of the solution with standard acid d. Measurement of the total volume of the
solution
b.
Measurement of the pH with a pH meter
c.
Determination of the boiling point of the
solution
e.
Measurement of the specific heat of the
solution
____
32. A rigid metal tank contains oxygen gas. Which of the following applies to the gas in the tank when
additional oxygen is added at constant temperature?
a. The volume of the gas increases.
d. The total number of gas molecules
remains the same.
b. The pressure of the gas decreases.
e. The average distance between the gas
molecules increases.
c. The average speed of the gas molecules
remains the same.
____
33. When hafnium metal is heated in an atmosphere of chlorine gas, the product of the reaction is found to
contain 62.2 percent Hf by mass and 37.4 percent Cl by mass. What is the empirical formula for this
compound?
a. HfCl
d. HfCl4
b. HfCl2
e. Hf2Cl3
c. HfCl3
____
34. Which of the following techniques is the most appropriate for the recovery of solid KNO 3 from an aqueous
solution of KNO3?
a. Paper chromatography
d. Electrolysis
b. Filtration
e. Evaporation to dryness
c. Titration
____
35. In the periodic table, as the atomic number increases from 11 to 17, what happens to the atomic radius?
a. It remains constant.
d. It decreases only.
b. It increases only.
e. It decreases, then increases.
c. It increases, then decreases.
____
36. Which of the following is a correct interpretation of the results of Rutherford's experiments in which gold
atoms were bombarded with alpha particles?
a. Atoms have equal numbers of positive and d. Neutrons and protons in atoms have
negative charges.
nearly equal mass.
b. Electrons in atoms are arranged in shells. e. The positive charge of an atom is
concentrated in a small region.
c. Neutrons are at the center of an atom.
____
37. Under which of the following sets of conditions could the most O 2(g) be dissolved in H2O(l)?
Pressure of O2(g) Above
Temperature of H2O(l) (oC)
H2O(l) (atm)
(A)
5.0
80
(B)
5.0
20
(C)
1.0
80
(D)
1.0
20
(E)
0.5
20
a.
b.
c.
____
A
B
C
d.
e.
D
E
38. Gases W and X react in a closed, rigid vessel to form gases Y and Z according to the equation above.
The intial pressure of W(g) is 1.20 atm and that of X(g) is 1.60 atm. No Y(g) or Z(g) is initially present. The
experiment is carried out at constant temperature. What is the partial pressure of Z(g) when the partial
pressure of W(g) has decreased to 1.0 atm?
a. 0.20 atm
d. 1.2 atm
b. 0.40 atm
e. 1.4 atm
c. 1.0 atm
____
39. 10 HI + 2 KMnO4 + 3 H2SO4  5 I2 + 2 MnSO4 + K2SO4 + 8 H2O
According to the balanced equation above, how many moles of HI would be necessary to produce 2.5
mol of I2, starting with 4.0 mol of KMnO4 and 3.0 mol of H2SO4 ?
a. 20.
d. 5.0
b. 10.
e. 2.5
c. 8.0
____
40. A yellow precipitate forms when 0.5 M NaI(aq) is added to a 0.5 M solution of which of the following ions?
a. Pb2+(aq)
d. SO42-(aq)
2+
b. Zn (aq)
e. OH-(aq)
2c. CrO4 (aq)
____
41. On a mountaintop, it is observed that water boils at 90 o, not at 100oC as at sea level. This phenomenon
occurs because on the mountaintop the
a. equilibrium water vapor pressure is higher d. water molecules have a higher average
due to the higher atmospheric pressure.
kinetic energy due to the lower
atmospheric pressure
b. equilibrium water vapor pressure is lower e. water contains a greater concentration of
due to the higher atmospheric pressure.
dissolved gases
c. equilibrium water vapor pressure equals
the atmospheric pressure at a lower
temperature
____
42. A 40.0 mL sample of 0.25 M KOH is added to 60.0 mL of 0.15 M Ba(OH)2. What is the molar
concentration of OH-(aq) in the resulting solution? (Assume that the volumes are additive.)
a. 0.10 M
d. 0.40 M
b. 0.19 M
e. 0.55 M
c. 0.28 M
____
43. NH4NO3(s)  N2O(g) + 2 H2O(g)
A 0.03 mol sample of NH4NO3(s) is placed in a 1 L evacuated flask, which is then sealed and heated. The
NH4NO3(s) decomposes completely according to the balanced equation above. The total pressure in the
flask measured at 400 K is closest to which of the following? (The value of the gas constant, R, is 0.082 L
atm mol-1 K-1.)
a. 3 atm
d. 0.1 atm
b. 1 atm
e. 0.03 atm
c. 0.5 atm
____
44. C2H4(g) + 3 O2(g)  2 CO2(g) + 2 H2O(g)
For the reaction of ethylene represented above, H is -1,323 kJ. What is the value of H if the
combustion produced liquid water H2O(g) ? (H for the phase change H2O(g)  H2O(l) is -44 kJ mol-1.)
a. -1,235 kJ
d. -1,367 kJ
b. -1,279 kJ
e. -1,411 kJ
c. -1,323 kJ
____
45. The graph above shows the results of a study of the reaction of X with a large excess of Y to yield Z. The
concentrations of X and Y were measured over a period of time. According to the results, which of the
following can be concluded about the rate law for the reaction under the conditions studied?
a. It is zero order in [X].
d. It is first order in [Y].
b. It is first order in [X].
e. The overall order of the reaction is 2.
c.
It is second order in [X].
____
46. Equal numbers of moles of He(g), Ar(g), and Ne(g) are placed in a glass vessel at room temperature. If
the vessel has a pinhole-sized leak, which of the following will be true regarding the relative values of the
partial pressures of the gases remaining in the vessel after some of the gas mixture has effused?
a. PHe < PNe<PAr
d. PAr<PHe< PNe
b. PHe<PAr<PNe
e. PAr=PHe= PNe
c. PNe<PAr<PHe
____
47. Which of the following compounds is NOT appreciably soluble in water but is soluble in dilute hydrochloric
acid?
a. Mg(OH)2(s)
d. (NH4)2SO4(s)
b. (NH4)2CO3(s)
e. Sr(NO3)2(s)
c. CuSO4(s)
____
48. In which of the following processes are covalent bonds broken?
a. I2(s)  I2(g)
d. C(diamond)  C(g)
b. CO2(s)  CO2(g)
e. Fe(s)  Fe(l)
c. NaCl(s)  NaCl(l)
____
49. What is the final concentration of barium ions, [Ba2+], in solution when 100. mL of 0.10 M BaCl2(aq) is
mixed with 100. mL of 0.050 M H2SO4(aq)?
a. 0.00 M
d. 0.075 M
b. 0.012 M
e. 0.10 M
c. 0.025 M
____
50. When 100 mL of 1.0 M Na3PO4 is mixed with 100 mL of 1.0 M AgNO3, a yellow precipitate forms when
[Ag+] becomes negligibly small. Which of the following is a correct listing of the ions remaining in solution
in order of increasing concentration?
a. [PO43-] < [NO3-] < [Na+]
d. [Na+] < [NO3-] < [PO43-]
3+
b. [PO4 ] < [Na ] < [NO3 ]
e. [Na+] < [PO43-] < [NO3-]
3+
c. [NO3 ] < [PO4 ] < [Na ]
____
51. In a qualitative analysis for the presence of Pb2+, Fe2+, and Cu2+ ions in aqueous solution, which of the
following will allow the separation of Pb2+ from the other ions at room temperature?
a. Adding dilute Na2S(aq) solution
d. Adding dilute NH3(aq) solution
b. Adding dilute HCl(aq) solution
e. Adding dilute HNO3(aq) solution
c. Adding dilute NaOH(aq) solution
____
52. After completing an experiment to determine gravimetrically the percentage of water in a hydrate, a
student reported a value of 38 percent. The correct value for the percentage of water in the hydrate is 51
percent. Which of the following is the most likely explanation for this difference?
a. Strong initial heating caused some of the d. The crucible was not heated to constant
hydrate sample to spatter out of the
mass before use.
crucible.
b. The dehydrated sample absorbed
e. Excess heating caused the dehydrated
moisture after heating.
sample to decompose.
c. The amount of the hydrate sample used
was too small.
____
53. The volume of distilled water that should be added to 10.0 mL of 6.00 M HCl(aq) in order to prepare a
0.500 M HCl(aq) solution is approximately
a. 50.0 mL
d. 110. mL
b. 60.0 mL
e. 120. mL
c. 100. mL
____
54. Which of the following gases deviates most from ideal behavior?
a. SO2
d. N2
b. Ne
e. H2
c. CH4
____
55. Which of the following pairs of liquids forms the solution that is most ideal (most closely follows Raoult's
law)?
a. C8H18(l) and H2O(l)
d. C6H14(l) and C8H18(l)
b. CH3CH2CH2OH(l) and H2O(l)
e. H2SO4(l) and H2O(l)
c. CH3CH2CH2OH(l) and C8H18(l)
____
56. The atom that contains exactly two unpaired electrons
a. S
d. Sb
b. Ca
e. Br
c. Ga
____
57. The atom that contains only one electron in the highest occupied energy sublevel
a. S
d. Sb
b. Ca
e. Br
c. Ga
____
58. The molecule with only one double bond
a. CO2
b. H2O
c. CH4
d.
e.
C2H4
PH3
59. The molecule with the largest dipole moment
a. CO2
b. H2O
c. CH4
d.
e.
C2H4
PH3
____
____
60. The molecule that has trigonal planar pyramidal geometry
a. CO2
d. C2H4
b. H2O
e. PH3
c. CH4
____
61. Is purple in aqueous solution
a. PbSO4
b. CuO
c. KMnO4
d.
e.
KCl
FeCl3
62. Is white and very soluble in water
a. PbSO4
b. CuO
c. KMnO4
d.
e.
KCl
FeCl3
____
____
63. Has an average atomic or molecular speed closest to that of N2 molecules at 0o and 1 atm
a. Ne
d. CO
b. Xe
e. NO
c. O2
____
64. Has the greatest density
a. Ne
b. Xe
c. O2
d.
e.
CO
NO
____
65. Has the greatest rate of effusion through a pinhole
a. Ne
d. CO
b. Xe
e. NO
c. O2
____
66. A precipitation reaction
a. H2SeO4(aq) + 2 Cl-(aq) + 2 H+(aq) 
H2SeO3(aq) + Cl2(g) + H2O(l)
b. S8(s) + 8 O2(g)  8 SO2
c. 3 Br2(aq) + 6 OH-(aq)  5 Br-(aq) + BrO3-
d.
Ca2+(aq) + SO42-(aq)  CaSO4(s)
e.
PtCl4(s) + 2 Cl-(aq)  PtCl62-
(aq) + 3 H2O(l)
____
67. A reaction in which the same reactant undergoes both oxidation and reduction
a. H2SeO4(aq) + 2 Cl-(aq) + 2 H+(aq) 
d. Ca2+(aq) + SO42-(aq)  CaSO4(s)
H2SeO3(aq) + Cl2(g) + H2O(l)
b. S8(s) + 8 O2(g)  8 SO2
e. PtCl4(s) + 2 Cl-(aq)  PtCl62c. 3 Br2(aq) + 6 OH-(aq)  5 Br-(aq) + BrO3(aq) + 3 H2O(l)
____
68. A combustion reaction
a. H2SeO4(aq) + 2 Cl-(aq) + 2 H+(aq) 
H2SeO3(aq) + Cl2(g) + H2O(l)
b. S8(s) + 8 O2(g)  8 SO2
c. 3 Br2(aq) + 6 OH-(aq)  5 Br-(aq) + BrO3(aq) + 3 H2O(l)
d.
Ca2+(aq) + SO42-(aq)  CaSO4(s)
e.
PtCl4(s) + 2 Cl-(aq)  PtCl62-
____
69. The substance is at its normal freezing point at time
a. t1
d. t4
b. t2
e. t5
c. t3
____
70. Which of the following best describes what happens to the substance between t4 and t5?
a. The molecules are leaving the liquid
d. The average intermolecular distance is
phase.
decreasing.
b. The solid and liquid phases coexist in
e. The temperature of the substance is
equilibrium.
increasing.
c. The vapor pressure of the substance is
decreasing.
____
71. In which of the following groups are the three species isoelectronic; i.e., have the same number of
electrons?
a. S2-, K+, Ca2+
d. Mg2+, Ca2+, Sr2+
2+
b. Sc, Ti, V
e. Cs, Ba2+, La3+
22c. O , S , Cl
____
72. The phase diagram for the pure substance X is shown above. The temperature of a sample of pure solid
X is slowly raised from 10oC to 100oC constant pressure of 0.5 atm. What is the expected behavior of the
substance?
a. It first melts to a liquid and then boils at
d. It sublimes to a vapor at an equilibrium
about 70oC.
temperature of about 20oC.
b. It first melts to a liquid and then boils at
e. It remains a solid until the temperature is
about 30oC.
greater than 100oC.
c. It melts to a liquid at a temperature of
about 20oC and remains a liquid until the
temperature is greater than 100oC.
____
73. In which of the following species does sulfur have the same oxidation number as it does in H2SO4?
a. H2SO3
d. S8
b. S2O32e. SO2Cl2
c. S2-
____
74. A flask contains 0.25 mole of SO2(g), 0.50 mole of CH4(g), and 0.50 mole of O2(g). The total pressure of
the gases in the flask is 800 mm Hg. What is the partial pressure of the SO2(g) in the flask?
a. 800 mm Hg
d. 200 mm Hg
b. 600 mm Hg
e. 160 mm Hg
c. 250 mm Hg
____
75. In the laboratory, H2(g) can be produced by adding which of the following to 1 M HCl(aq)?
I. 1 M NH3(aq)
II. Zn(s)
III. NaHCO3(s)
a. I only
d. I and II only
b. II only
e. I, II, and III
c. III only
____
76. A compound contains 1.10 mol of K, 0.55 mol of Te, and 1.65 mol of O. What is the simplest formula of
this compound?
a. KTeO
d. K2TeO6
b. KTe2O
e. K4TeO6
c. K2TeO3
____
77. 3 C2H2(g)  C6H6(g)
What is the standard enthalpy change, Ho, for the reaction represented above?
Hof of C2H2(g) is 230 kJ mol-1; Hof of C6H6(g) is 83 kJ mol-1.)
a. -607 kJ
d. +19 kJ
b. -147 kJ
e. +773 kJ
c. -19 kJ
____
78. Approximately what mass of CuSO4•5H2O (250 g mol-1) is required to prepare 250 mL of 0.10 M
copper(II) sulfate solution?
a. 4.0 g
d. 85 g
b. 6.2 g
e. 140 g
c. 34 g
____
79. 2 NO(g) + O2(g) 2 NO2(g)
A possible mechanism for the overall reaction represented above is the following.
(1) NO(g) + NO(g)  N2O2(g)
slow
(2) N2O2(g) + O2(g)  2 NO2(g)
fast
Which of the following rate expressions agrees best with this possible mechanism?
a. Rate = k[NO]2
d. Rate = k[NO]2[O2]
b.
e. Rate = k[N2O2][O2]
Rate = k
c.
Rate = k
____
80. Of the following compounds, which is the most ionic?
a. SiCl4
d. Cl2O
b. BrCl
e. CaCl2
c. PCl3
____
81. The best explanation for the fact that diamond is extremely hard is that diamond crystals
a. are made up of atoms that are intrinsically d. are formed under extreme conditions of
hard because of their electronic structures
temperature and pressure
b.
consist of positive and negative ions that
are strongly attracted to each other
c.
are giant molecules in which each atom
forms strong covalent bonds with all of its
neighboring atoms
e.
contain orbitals or bands of delocalized
electrons that belong not to single atoms
but to each crystal as a whole
____
82. CS2(l) + 3 O2(g)  CO2(g) + 2 SO2(g)
What volume of O2(g) is required to resact with excess CS2(l) to produce 4.0 L of CO2(g)? (Assume all
gases are measured at 0oC and 1 atm.)
a. 12 L
d. 2 X 22.4 L
b. 22.4 L
e. 3 X 22.4 L
c.
X 22.4 L
____
83. Which of the following oxides is a gas at 25oC and 1 atm?
a. Rb2O
d. SiO2
b. N2O
e. La2O3
c. Na2O2
____
84. A solution is made by dissolving a nonvolatile solute in a pure solvent. Compared to the pure solvent, the
solution
a. has a higher boiling point
d. has a higher freezing point
b. has a higher vapor pressure
e. is more nearly ideal
c. has the same vapor pressure
____
85. A sample of a solution of an unknown was treated with dilute hydrochloric acid. The white precipitate
formed was filtered and washed with hot water. A few drops of potassium iodide solution were added to
the hot water filtrate and a bright yellow precipitate was produced. The white precipitate remaining on the
filter paper was readily soluble in ammonia solution. What two ions could have been present in the
unknown?
a. Ag+(aq) and Hg22+(aq)
d. Ba2+(aq) and Hg22+(aq)
+
2+
b. Ag (aq) and Pb (aq)
e. Ba2+(aq) and Pb2+(aq)
2+
+
c. Ba (aq) and Ag (aq)
____
86. A 0.10 M aqueous solution of sodium sulfate, Na2SO4, is a better conductor of electricity than a 0.10 M
aqueous solution of sodium chloride, NaCl. Which of the following best explains this observation?
a. Na2SO4 is more solutble in water than
d. More moles of ions are present in a given
volume of 0.10 M Na2SO4 than in the
NaCl is.
same volume of 0.10 M NaCl.
b. Na2SO4 has a higher molar mass than
e. The degree of dissociation of Na2SO4 in
NaCl has.
solution is significantly greater than that of
NaCl.
c. To prepare a given volume of 0.10 M
solution, the mass of Na2SO4 needed is
more than twice the mass of NaCl needed.
____
87. On the basis of the solubility curves shown above, the greatest percentage of which compound can be
recovered by cooling a saturated solution of that compound from 90o to 30o?
a. NaCl
d. K2SO4
b. KNO3
e. Ce2(SO4)3
c. K2CrO4
____
88. An excess of Mg(s) is added to 100. mL of 0.400 M HCl. At 0oC and 1 atm pressure, what volume of H2
gas can be obtained?
a. 22.4 mL
d. 448 mL
b. 44.8 mL
e. 896 mL
c. 224 mL
____
89. The atomic mass of copper is 63.55. Given that there are only two naturally occurring isotopes of copper,
63Cu and 65Cu, the natural abundance of the 65Cu isotope must be approximately
a. 90%
d. 25%
b. 70%
e. 10%
c. 50%
____
90. Which of the following properties generally decreases across the periodic table from sodium to chlorine?
a. First ionization energy
d. Maximum value of oxidation number
b. Atomic mass
e. Atomic radius
c. Electrongeativity
____
91. What is the mole fraction of ethanol, C2H5OH, in an aqueous solution that is 46% ethanol by mass? (The
molar mass of C2H5OH is 46 g; the molar mass of H2O is 18 g.)
a. 0.25
d. 0.67
b. 0.46
e. 0.75
c. 0.54
____
92. The effective nuclear charge experienced by the outermost electron of Na is different than the effective
nuclear charge experienced by the outermost electron of Ne. This difference best accounts for which of
the following?
a. Na has a greater density at standard
d. Na has a higher neutron-to-proton ratio
conditions than Ne.
than Ne.
b. Na has a lower first ionization energy than e. Na has a fewer naturally occuring isotopes
Ne.
than Ne.
c. Na has a higher melting point than Ne.
____
93. Which of the following is a correct statement about reaction order?
a. Reaction order can only be a whole
d. Reaction order increases with increasing
number.
temperature.
b. Reaction order can be determined only
e. A second-order reaction must involve at
from the coefficients of the balanced
least two different compounds as
equation for the reaction.
reactants.
c. Reaction order can be determined only by
experiment.
____
94. Sodium chloride is LEAST soluble in which of the following liquids?
a. H2O
d. CH3OH
b. CCl4
e. CH3COOH
c. HF
____
95. ...Cr2O72-(aq) + ...H2S(g) + ...H+(aq)  ...Cr3+(aq) + ...S(s) + ...H2O(l)
When the equation above is correctly balanced and all coefficients are reduced to lowest whole-number
terms, the coefficient for H+(aq) is
a. 2
d. 8
b. 4
e. 14
c. 6
____
96. Which of the following represents acceptable laboratory practice?
a. Placing a hot object on a balance pan
d. Using 10 mL of standard strength
b.
c.
Using distilled water for the final rinse of a e.
buret before filling it with standardized
solution
Adding a weighed quantity of solid acid to
a titration flask wet with distilled water
phenolphthalein indicator solution for
titration of 25 mL of acid solution
Diluting a solution in a volumetric flask to
its final concentration with hot water
____
97. 3 Cu(s) + 8 H+(aq) + 2 NO3-(aq)  3 Cu2+(aq) + 2 NO(g) + 4 H2O(l)
True statements about the reaction represented above include which of the following?
I. Cu(s) acts as an oxidizing agent.
II. The oxidation state of nitrogen changes from +5 to +2.
III. Hydrogen ions are oxidized to form H2O(l).
a. I only
d. I and II
b. II only
e. II and III
c. III only
____
98. Propane gas, C3H8, burns in excess oxygen gas. When the equation for this reaction is correctly balanced
and all coefficients are reduced to their lowest whole-number term, the coefficient for O2 is
a. 4
d. 10
b. 5
e. 22
c. 7
____
99. According to the VSEPR model, the progressive decrease in the bond angles in the series of molecules
CH4, NH3, and H2O is best accounted for by the
a. increasing strength of the bonds
d. increasing the number of unshared pairs
of electrons
b. decreasing size of the central atom
e. decreasing repulsion between hydrogen
atoms
c. increasing electronegativity of the central
atom
100. 0
1
2
3
4
5
6
7
...
10
...
20
Tim
e (days)
% Reactant 100
79
63
50
40
31
25
20
10
1
remaining
A reaction was observed for 20 days and the percentage of the reactant remaining after each day was
recorded in the table above. Which of the following best describes the order and the half-life of the
reaction?
a. Reaction Order: First, Half-life (days): 3
d. Reaction Order: Second, Half-life (days): 6
b. Reaction Order: First, Half-life (days): 10 e. Reaction Order: Second, Half-life (days):
10
c. Reaction Order: Second, Half-life (days): 3
____
____ 101. The boiling points of the elements helium, neon, argon, krypton, and xenon increase in that order. Which
of the following statements accounts for this increase?
a. The London (dispersion) forces increase. d. The chemical reactivity increases.
b. The hydrogen bonding increases.
e. The number of nearest neighbors
increases.
c. The dipole-dipole forces increase.
____ 102. Rate = k[M][N]2
The rate of a certain chemical reaction between substances M and N obeys the rate law above. The
reaction is first studied with [M] and [N] each 1 X 10-3 molar. If a new experiment is conducted with [M]
and [N] each 2 X 10-3 molar, the reaction rate will increase by a factor of
a. 2
d. 8
b. 4
e. 16
c. 6
____ 103. 2 N2H4(g) + N2O4(g)  3 N2(g) + 4 H2O(g)
When 8.0 g of N2H4 (32 g mol-1) and 92 g of N2O4 (92 g mol-1) are mixed together and react according to
the equation above, what is the maximum mass of H2O that can be produced?
a. 9.0 g
d. 72 g
b. 18 g
e. 144 g
c. 36 g
____ 104. All of the halogens in their elemental form at 25oC and 1 atm are
a. conductors of electricity
d. colorless
b. diatomic molecules
e. gases
c. odorless
____ 105. 2 H2O(l) + 4 MnO4-(aq) + 3 ClO2-(aq)  4 MnO2(s) + 3 ClO4-(aq) + 4 OH-(aq)
According to the balanced equation above, how many moles of ClO 2-(aq) are needed to react completely
with 20. mL of 0.20 M KMnO4 solution?
a. 0.0030 mol
d. 0.013 mol
b. 0.0053 mol
e. 0.030 mol
c. 0.0075 mol
____ 106. Which of the following structural formulas represents an isomer of the compound that has the structural
formula represented above?
a.
d.
b.
c.
e.
____ 107. Which of the following substances is LEAST soluble in water?
a. (NH4)2SO4
d. Zn(NO3)2
b. KMnO4
e. Na3PO4
c. BaCO3
____ 108. A 2 L container will hold about 4 g of which of the following gases at 0oC and 1 atm?
a. SO2
d. C4H8
b. N2
e. NH3
c. CO2
____ 109. Which of the following describes the changes in forces of attraction that occur as H 2O changes phase
from a liquid to a vapor?
a. H-O bonds break as H-H and O-O bonds d. Ionic bonds between H+ ions and OH- ions
form.
are broken.
b. Hydrogen bonds between H2O molecules e. Covalent bonds between H+ ions and H2O
are broken.
molecules become more effective.
c. Covalent bonds between H2O molecules
are broken.
____ 110. Liquid naphthalene at 95oC was cooled to 30oC, as represented in the cooling curve above. From which
section of the curve can the melting point of naphthalene be determined?
a. A
d. D
b. B
e. E
c. C
____ 111. If 200. mL of 0.60 M MgCl2(aq) is added to 400. mL of distilled water, what is the concentration of
Mg2+(aq) in the resulting solution? (Assume volumes are additive.)
a. 0.20 M
d. 0.60 M
b. 0.30 M
e. 1.2 M
c. 0.40 M
____ 112. Of the following pure substances, which has the highest melting point?
a. S8
d. SO2
b. I2
e. C6H6
c. SiO2
____ 113. A colorless solution is divided into three samples. The following tests were performed on samples of the
solution.
Sample
Test
Observation
Add H+(aq)
1
No change
Add NH3(aq)
2
No change
Add SO42-(aq)
3
No change
Which of the following ions could be present in the solution at a concentration of 0.10 M ?
a. Ni2+(aq)
d. Na+(aq)
3+
b. Al (aq)
e. CO32-(aq)
2+
c. Ba (aq)
____ 114. X(s)  X(l)
Which of the following is true for any substance undergoing the process represented above at its normal
melting point?
a. S < 0
d. TS = 0
b. H = 0
e. H = TS
c. H = TDS
____ 115. Forms monatomic ions with 2¯ charge in solutions
a. F
d. Ar
b. S
e. Mn
c. Mg
____ 116. Forms a compound having the formula KXO4
a. F
b. S
c. Mg
d.
e.
Ar
Mn
____ 117. Forms oxides that are common air pollutants and that yield acidic solution in water
a. F
d. Ar
b. S
e. Mn
c. Mg
____ 118. Is a good oxidizing agent
a. Hydrofluoric acid
b. Carbon dioxide
c. Aluminum hydroxide
d.
e.
Ammonia
Hydrogen peroxide
____ 119. Is used to etch glass chemically
a. Hydrofluoric acid
b. Carbon dioxide
c. Aluminum hydroxide
d.
e.
Ammonia
Hydrogen peroxide
____ 120. Is used extensively for the production of fertilizers
a. Hydrofluoric acid
d.
b. Carbon dioxide
e.
c. Aluminum hydroxide
Ammonia
Hydrogen peroxide
____ 121. Solid ethyl alcohol, C2H5OH
a. A network solid with covalent bonding
d.
b. A molecular solid with zero dipole moment e.
c. A molecular solid with hydrogen bonding
An ionic solid
A metallic solid
____ 122. Silicon dioxide, SiO2
a. A network solid with covalent bonding
d.
b. A molecular solid with zero dipole moment e.
c. A molecular solid with hydrogen bonding
An ionic solid
A metallic solid
____ 123. Assume that you have an "unknown" consisting of an aqueous solution of a salt that contains one of the
ions listed below. Which ions must be absent on the basis of each of the following observations of the
"unknown"?
The solution is colorless
a. CO32¯
b. Cr2O72¯
c. NH4+
d.
e.
Ba2+
Al3+
____ 124. Assume that you have an "unknown" consisting of an aqueous solution of a salt that contains one of the
ions listed below. Which ions must be absent on the basis of each of the following observations of the
"unknown"?
The solution gives no apparent reaction with dilute hydrochloric acid.
a. CO32¯
d. Ba2+
2
b. Cr2O7 ¯
e. Al3+
c. NH4+
____ 125. Assume that you have an "unknown" consisting of an aqueous solution of a salt that contains one of the
ions listed below. Which ions must be absent on the basis of each of the following observations of the
"unknown"?
No odor can be detected when a sample of the solution is added drop by drop to a warm solution of
sodium hydroxide.
a.
b.
c.
CO32¯
Cr2O72¯
NH4+
d.
e.
Ba2+
Al3+
____ 126. Assume that you have an "unknown" consisting of an aqueous solution of a salt that contains one of the
ions listed below. Which ions must be absent on the basis of each of the following observations of the
"unknown"?
No precipitate is formed when a dilute solution of H2SO4 is added to a sample of the solution.
a. CO32¯
d. Ba2+
b. Cr2O72¯
e. Al3+
c. NH4+
____ 127.
Hydrogen Halide
HF
HCl
HBr
HI
Normal Boiling Point, °C
+19
- 85
- 67
- 35
The liquefied hydrogen halides have the normal boiling points given above. The relatively high boiling
point of HF can be correctly explained by which of the following?
a.
b.
HF gas is more ideal.
HF is the strongest acid.
c.
HF molecules have a smaller dipole
moment.
d.
e.
HF is much less soluble in water.
HF molecules tend to form hydrogen
bonds.
____ 128. Which of the following represents a pair of isotopes?
a. I. Atomic Number: 6, Mass Number: 14
d. I. Atomic Number: 7, Mass Number: 13
II. Atomic Number: 7, Mass Number, 14
II. Atomic Number: 7, Mass Number: 14
b. I. Atomic Number: 6, Mass Number: 7
e. I. Atomic Number: 8, Mass Number: 16
II. Atomic Number: 14, Mass Number: 14
II. Atomic Number: 16, Mass Number: 20
c. I. Atomic Number: 6, Mass Number: 14
II. Atomic Number: 14, Mass Number: 28
____ 129. .....Mg(s) + .....NO3¯(aq) +.....H+(aq) --->......Mg2+(aq) + ....NH4+(aq) + ....H2O(l)
When the skeleton equation above is balanced and all coefficients reduced to their lowest whole-number
terms. what is the coeficient for H+ ?
a.
b.
c.
4
6
8
d.
e.
9
10
____ 130. When a sample of oxygen gas in a closed container of constant volume is heated until its absolute
temperature is doubled, which of the following is also doubled?
a.
b.
c.
d.
The density of the gas
e.
The pressure of the gas
The average velocity of the gas molecules
The number of molecules per cm 3
The potential energy of the molecules
____ 131. 1s2 2s22p6 3s23p3
Atoms of an element, X, have the electronic configuration shown above. The compound most likely
formed with magnesium, Mg, is
a.
b.
c.
MgX
Mg2X
MgX2
d.
e.
MgX3
Mg3X2
____ 132. The density of an unknown gas is 4.20 grams per liter at 3.00 atmospheres pressure and 127 °C. What is
the molecular weight of this gas? (R = 0.0821 liter-atm / mole-K)
a. 14.6
d. 94.1
b. 46.0
e. 138
c. 88.0
____ 133. The formula for potassium hexacyanoferrate(II) is
a. K4[Fe(CN)6]
d.
b. K3[Fe(CN)6]
e.
c. K2[Pt(CN)4]
K2[Pt(CN)6]
KCN
____ 134.
H3AsO4 + 3I¯ + 2 H3O+ ---> H3AsO3 + I3¯ + H2O
The oxidation of iodide ions by arsenic acid in acidic aqueous solution occurs according to the
stoichiometry shown above. The experimental rate law of the reaction is:
Rate = k [H3AsO4] [I¯] [H3O+]
What is the order of the reaction with respect to I¯?
a. 1
d. 5
b. 2
e. 6
c. 3
____ 135.
H3AsO4 + 3I¯ + 2 H3O+ ---> H3AsO3 + I3¯ + H2O
The oxidation of iodide ions by arsenic acid in acidic aqueous solution occurs according to the
stoichiometry shown above. The experimental rate law of the reaction is:
Rate = k [H3AsO4] [I¯] [H3O+]
According to the rate law for the reaction, an increase in the concentration of hydronium ion has what
effect on this reaction?
a. The rate of reaction increases.
d. The value of the equilibrium constant
decreases.
b. The rate of reaction decreases.
e. Neither the rate nor the value of the
equilibrium constant is changed.
c. The value of the equilibrium constant
increases.
____ 136. The critical temperature of a substance is the
a. temperature at which the vapor pressure
of the liquid is equal to the external
pressure
b. temperature at which the vapor pressure
of the liquid is equal to 760 mm Hg
c.
d.
Temperature at which liquid and vapor
phases are in equilibrium at I atmosphere
e.
lowest temperature above which a
substance cnnot be liquified at any applied
pressure
temperature at which the solid, liquid, and
vapor phases are all in equilibrium
____ 137. 2 A(g) + B(g) <===> 2 C(g)
When the concentration of substance B in the reaction above is doubled, all other factors being held
constant, it is found that the rate of the reaction remains unchanged. The most probable explanation for
this observation is that
a.
the order of the reaction with respect to
substance B is 1
d.
b.
substance B is not involved in any of the
steps in the mechanism of the reaction
e.
c.
substance B is not involved in the ratedetermining step of the mechanism, but is
involved in subsequent steps
substance B is probably a catalyst, and as
such, its effect on the rate of the reaction
does not depend on its concentration
the reactant with the smallest coefficient in
the balanced equation generally has little
or no effecton the rate of the reaction
____ 138. A 0.1-molar solution of which of the following ions is orange?
a. Fe(H2O)42+
d. Zn(NH3)42+
2+
b. Cu(NH3)4
e. Cr2O72¯
2
c. Zn(OH)4 ¯
____ 139. The net ionic equation for the reaction between silver carbonate and hydrochloric acid is
a. Ag2CO3(s) + 2 H+ + 2 Cl¯ ---> 2 AgCl(s) + d. Ag+ + Cl¯ ---> AgCl(s)
H2O + CO(g)
b. 2 Ag+ + CO32¯ + 2 H+ + 2 Cl¯ ---> 2
e. Ag2CO3(s) + 2 H+ ---> 2Ag+ + H2CO3
AgCl(s) + H2O + CO2(g)
c. CO32¯ + 2 H+ ---> H2O + CO2(g)
____ 140. ...CrO2¯ + ...OH¯ ---> ... CrO42¯ + ... H2O + ... e¯ When the equation for the half-reaction above is
balanced, what is the ratio of the coefficients OH¯ / CrO 2¯ ?
a. 1:1
d. 4:1
b. 2:1
e. 5:1
c. 3:1
____ 141. The addition of an oxidizing agent such as chlorine water to a clear solution of an unknown compound
results in the appearance of a brown color. When this solution is shaken with the organic solvent,
methylene dichloride, the organic solvent layer turns purple. The unknown compound probably contains
a. K+
d. I¯
b. Br¯
e. Co2+
c. NO3¯
____ 142. The molality of the glucose in a 1.0-molar glucose solution can be obtained by using which of the
following?
a.
b.
c.
Volume of the solution
Temperature of the solution
Solubility of glucose in water
d.
e.
Degree of dissociation of glucose
Density of the solution
____ 143. Equal masses of three different ideal gases, X, Y, and Z, are mixed in a sealed rigid container. If the
temperature of the system remains constant, which of the following statements about the partial pressure
of gas X is correct?
a. It is equal to 1/3 the total pressure
d. It depends on the average distance
traveled between molecular collisions.
b. It depends on the intermolecular forces of e. It can be calculated with knowledge only
attraction between molecules of X, Y, and
of the volume of the container.
Z.
c. It depends on the relative molecular
masses of X, Y, and Z.
____ 144. The geometry of the SO3 molecule is best described as
a. trigonal planar
d. bent
b. trigonal pyramidal
e. tetrahedral
c. square pyramidal
____ 145. Which of the following molecules has the shortest bond length?
a. N2
d. Br2
b. O2
e. I2
c. Cl2
____ 146. Metallic copper is heated strongly with concentrated sulfuric acid. The products of this reaction are
a. CuSO4(s) and H2(g) only
d. CuSO4(s), H2(g), and SO2(g)
b. Cu2+, SO2(g), and H2O
e. Cu2+, SO3(g), and H2O
2+
c. Cu , H2(g), and H2O
____ 147. The elements in which of the following have most nearly the same atomic radius?
a. Be, B, C, N
d. C, P, Se, I
b. Ne, Ar, Kr, Xe
e. Cr, Mn, Fe, Co
c. Mg, Ca, Sr, Ba
____ 148. What number of moles of O2 is needed to produce 14.2 grams of P4O10 from P? (Molecular weight P4O10
= 284)
a. 0.0500 mole
d. 0.250 mole
b. 0.0625 mole
e. 0.500 mole
c. 0.125 mole
____ 149. The alkenes are compounds of carbon and hydrogen with the general formula CnH2n. If 0.561 gram of any
alkene is burned in excess oxygen, what number of moles of H 2O is formed?
a.
b.
c.
0.0400 mole
0.0600 mole
0.0800 mole
d.
e.
0.400 mole
0.800 mole
____ 150. CH4(g) + 2 O2(g) ---> CO2(g) + 2 H2O(l); D = - 889.1 kJ
DHf° H2O(l) = - 285.8 kJ / mole
DHf° CO2(g) = - 393.3 kJ / mole
What is the standard heat of formation of methane, DHf° CH4(g), as calculated from the data above?
a.
b.
c.
-210.0 kJ/mole
-107.5 kJ/mole
-75.8 kJ/mole
d.
e.
75.8 kJ/mole
210.0 kJ/mole
____ 151. Two flexible containers for gases are at the same temperature and pressure. One holds 0.50 gram of
hydrogen and the other holds 8.0 grams of oxygen. Which of the following statements regarding these
gas samples is FALSE?
a. The volume of the hydrogen container is
d. The average kinetic energy of the
the same as the volume of the oxygen
hydrogen molecules is the same as the
container.
average kinetic energy of the oxygen
molecules.
b. The number of molecules in the hydrogen e. The average speed of the hydrogen
container is the same as the number of
molecules is the same as the average
molecules in the oxygen container.
speed of the oxygen molecules.
c. The density of the hydrogen sample is
less than that of the oxygen sample.
____ 152. Pi bonding occurs in each of the following species EXCEPT
a. CO2
d. C6H6
b. C2H4
e. CH4
c. CN¯
____ 153. 3 Ag(s) + 4 HNO3 <===> 3 AgNO3 + NO(g) + 2 H2O
The reaction of silver metal and dilute nitric acid proceeds according to the equation above. If 0.10 mole
of powdered silver is added to 10. milliliters of 6.0-molar nitric acid, the number of moles of NO gas that
can be formed is
a.
b.
c.
0.015 mole
0.020 mole
0.030 mole
d.
e.
0.045 mole
0.090 mole
____ 154. A cube of ice is added to some hot water in a rigid, insulated container, which is then sealed. There is no
heat exchange with the surroundings. What has happened to the total energy and the total entropy when
the system reaches equilibrium?
a. Energy: constant, Entropy: constant
d. Energy: decreases, Entropy: increases
b. Energy: constant, Entropy: decreases
e. Energy: increases, Entropy: decreases
c. Energy: constant, Entropy: increases
____ 155. Which of the following represents the ground state electron configuration for the Mn 3+ ion? (Atomic
number Mn = 25)
a. 1s2 2s22p6 3s23p63d4
d. 1s2 2s22p6 3s23p63d8 4s2
2
2
6
2
6
5
2
b. 1s 2s 2p 3s 3p 3d 4s
e. 1s2 2s22p6 3s23p63d3 4s1
2
2
6
2
6
2
2
c. 1s 2s 2p 3s 3p 3d 4s
____ 156. Which of the following has a zero dipole moment?
a. HCN
d. NO2
b. NH3
e. PF5
c. SO2
____ 157. When a solution of potassium dichromate is added to an acidified solution of iron(II) sulfate, the products
of the reaction are
a. FeCr2O7(s) and H2O
d. Fe3+, Cr3+, and H2O
b. FeCrO4(s) and H2O
e. Fe2(SO4)3(s), Cr3+ and H2O
c. Fe3+, CrO42¯, and H2O
____ 158. A student pipetted five 25.00-milliliter samples of hydrochloric acid and transferred each sample to an
Erlenmeyer flask, diluted it with distilled water, and added a few drops of phenolphthalein to each. Each
sample was then titrated with a sodium hydroxide solution to the appearance of the first permanent faint
pink color. The following results were obtained.
Volumes of NaOH Solution
First Sample..................35.22 mLSecond Sample..............36.14 mLThird Sample.................36.13
mLFourth Sample..............36.15 mLFifth Sample..................36.12 mL
Which of the following is the most probable explanation for the variation in the student's results?
a.
b.
c.
d.
The burette was not rinsed with NaOH
solution.
e.
The student misread a 5 for a 6 on the
burette when the first sample was titrated.
A different amount of water was added to
the first sample.
The pipette was not rinsed with the HCI
solution.
The student added too little indicator to
the first sample.
____ 159. The net ionic equation for the reaction that occurs during the titration of nitrous aicd with sodium
hydroxide is
a. HNO2 + Na+ + OH¯ ---> NaNO2 + H2O
d. HNO2 + H2O ---> NO2¯ + H3O+
b. HNO2 + NaOH ---> Na+ + NO2¯ + H2O
e. HNO2 + OH¯ ---> NO2¯ + H2O
c. H+ + OH¯ --->H2O
____ 160. Which of the following species CANNOT function as an oxidizing agent?
a. Cr2O72¯
d. S
b. MnO4¯
e. I¯
c. NO3¯
____ 161. Ca, V, Co, Zn, As
Gaseous atoms of which of the elements above are paramagnetic?
a.
b.
c.
Ca and As only
Zn and As only
Ca, V, and Co only
d.
e.
V, Co, and As only
V, Co, and Zn only
____ 162. A student wishes to prepare 2.00 liters of 0.100-molar KIO3 (molecular weight 214). The proper procedure
is to weigh out
a. 42.8 grams of KIO3 and add 2.00
d. 42.8 grams of KIO3 and add 2.00 liters of
kilograms of H2O
H2O
b. 42.8 grams of KIO3 and add H2O until the e. 21.4 grams fo KIO3 and add 2.00 liters of
final homogeneous solution has a volume
H2O
of 2.00 liters
c. 21.4 grams of KIO3 and add H2O until the
final homogeneous solution has a volume
of 2.00 liters
____ 163. A 20.0-milliliter sample of 0.200-molar K2CO3 solution is added to 30.0 milliliters of 0.400-molar Ba(NO3)2
solution. Barium carbonate precipitates. The concentration of barium ion, Ba2+, in solution after reaction is
a. 0.150 M
d. 0.240 M
b. 0.160 M
e. 0.267 M
c. 0.200 M
____ 164. What is the mole fraction of ethanol, C2H5OH, in an aqueous solution in which the ethanol concentration
is 4.6 molal?
a. 0.0046
d. 0.20
b. 0.076
e. 0.72
c. 0.083
____ 165. One of the outermost electrons in a strontium atom in the ground state can be described by which of the
following sets of four quantum numbers?
a. 5, 2, 0, 1/2
d. 5, 0, 1, 1/2
b. 5, 1, 1, 1/2
e. 5, 0, 0, 1/2
c. 5, 1, 0, 1/2
____ 166. A compound is heated to produce a gas whose molecular weight is to be determined. The gas is
collected by displacing water in a water-filled flask inverted in a trough of water. Which of the following is
necessary to calculate the molecular weight of the gas, but does NOT need to be measured during the
experiment?
a. Mass of the compound used in the
d. Barometric pressure
experiment
b. Temperature of the water in the trough
e. Volume of water displaced from the flask
c. Vapor pressure of the water
____ 167. A 27.0-gram sample of an unknown hydrocarbon was burned in excess oxygen to form 88.0 grams of
carbon dioxide and 27.0 grams of water. What is a possible molecular formula of the hydrocarbon?
a. CH4
d. C4H6
b. C2H2
e. C4H10
c. C4H3
____ 168. When the actual gas volume is greater then the volume predicted by the ideal gas law, the explanation
lies in the fact that the ideal gas law does NOT include a factor for molecular.
a. volume
d. attractions
b. mass
e. shape
c. velocity
____ 169. 5 Fe2+ + MnO4¯ + 8 H+ <===> 5 Fe3+ + Mn2+ + 4 H2O
In a titration experiment based on the equation above, 25.0 milliliters of an acidified Fe 2+ solution requires
14.0 milliliters of standard 0.050-molar MnO4¯ solution to reach the equivalence point. The concentration
of Fe2+ in the original solution is
a.
b.
c.
0.0010 M
0.0056 M
0.028 M
d.
e.
0.090 M
0.14 M
____ 170. For which of the following molecules are resonance structures necessary to describe the bonding
satisfactorily?
a. H2S
d. OF2
b. SO2
e. PF3
c. CO2
____ 171. What is the net ionic equation for the reaction that occurs when aqueous copper(II) sulfate is added to
excess 6-molar ammonia?
a. Cu2+ + SO42¯ + 2 NH4+ + 2 OH¯ --->
d. Cu2+ + 4 NH3 --> Cu(NH3)42+
(NH4)2SO4 + Cu(OH)2
b. Cu2+ + 4 NH3 + 4 H2O --> Cu(OH)42¯ + 4
e. Cu2+ + 2 NH3 + H2O --> CuO + 2 NH4+
NH4+
c.
Cu2+ + 2 NH3 + 2 H2O --> Cu(OH)2 + 2
NH4+
____ 172. NH3(g) + 2 CH4(g) + 5/2 O2(g) <===> H2NCH2COOH(s) + 3 H2O(l)
At constant temperature, DH, the change in enthalpy for the reaction above is approximately equal to
a.
b.
c.
DE - (11/2)RT
DE - (7/2)RT
DE + RT
d.
e.
DE + (7/2)RT
DE (11/2)RT
____ 173. Which of the following aqueous solutions has the highest boiling point?
a. 0.10 M potassium sulfate, K2SO4
d. 0.10 M magnesium sulfalte, MgSO4
b. 0.10 M hydrochloric acid, HCl
e. 0.20 M sucrose, C12H22O11
c. 0.10 M ammonium nitrate, NH4NO3
____ 174. A sample of 9.00 grams of aluminum metal is added to an excess of hydrochloric acid. The volume of
hydrogen gas produced at standard temperature and pressure is
a. 22.4 liters
d. 5.60 liters
b. 11.2 liters
e. 3.74 liters
c. 7.46 liters
____ 175. Which element exhibits the greatest number of different oxidation states?
a. O
d. Mg
b. La
e. N
c. Rb
____ 176. Which of the elements above has the smallest ionic radius for its most commonly found ion?
a. O
d. Mg
b. La
e. N
c. Rb
____ 177. An impossible electronic configuration
a. 1s2 2s22p5 3s23p5
b. 1s2 2s22p6 3s23p6
c. 1s2 2s22p62d10 3s23p6
d.
e.
1s2 2s22p6 3s23p63d5
1s2 2s22p6 3s23p63d3 4s2
____ 178. The ground-state configuration for the atoms of a transition element
a. 1s2 2s22p5 3s23p5
d. 1s2 2s22p6 3s23p63d5
b. 1s2 2s22p6 3s23p6
e. 1s2 2s22p6 3s23p63d3 4s2
2
2
6
10
2
6
c. 1s 2s 2p 2d 3s 3p
____ 179. The ground-state configuration of a negative ion of a halogen
a. 1s2 2s22p5 3s23p5
d. 1s2 2s22p6 3s23p63d5
b. 1s2 2s22p6 3s23p6
e. 1s2 2s22p6 3s23p63d34s2
c. 1s2 2s22p62d103s23p6
____ 180. The ground-state configuration of a common ion of an alkaline earth element
a. 1s2 2s22p5 3s23p5
d. 1s2 2s22p6 3s23p63d5
2
2
6
2
6
b. 1s 2s 2p 3s 3p
e. 1s2 2s22p6 3s23p63d3 4s2
2
2
6
10
2
6
c. 1s 2s 2p 2d 3s 3p
____ 181. Is used to explain why iodine molecules are held together in the solid state
a. hydrogen bonding
d. resonance
b. hybridization
e. van der Waals forces (London dispersion
forces)
c. ionic bonding
____ 182. Is used to explain why the boiling point of HF is greater than the boiling point of HBr
a. hydrogen bonding
d. resonance
b. hybridization
e. van der Waals forces (London dispersion
forces)
c. ionic bonding
____ 183. Is used to explain the fact that the four bonds in methane are equivalent
a.
b.
hydrogen bonding
hybridization
c.
ionic bonding
d.
e.
resonance
van der Waals forces (London dispersion
forces)
____ 184. Is used to explain the fact that the carbon-to-carbon bonds in benzene, C6H6, are identical
a. hydrogen bonding
d. resonance
b. hybridization
e. van der Waals forces (London dispersion
forces)
c. ionic bonding
____ 185. The weight of H2SO4 (molecular weight 98.1) in 50.0 milliliters of a 6.00-molar solution is
a.
b.
c.
3.10 grams
12.0 grams
29.4 grams
d.
e.
294 grams
300. grams
____ 186. A gaseous mixture containing 7.0 moles of nitrogen, 2.5 moles of oxygen, and 0.50 mole of helium exerts a total
pressure of 0.90 atmosphere. What is the partial pressure of the nitrogen?
a. 0.13 atm
d. 0.90 atm
b. 0.27 atm
e. 6.3 atm
c. 0.63 atm
____ 187. The Lewis dot structure of which of the following molecules shows only one unshared pair of valence electron?
a. Cl2
d. CCl4
b. N2
e. H2O2
c. NH3
____ 188. 6 I¯ + 2 MnO4¯ + 4 H2O(l) ---> 3 I2(s) + 2 MnO2(s) + OH¯
Which of the following statements regarding the reaction represented by the equation above is correct?
a.
Iodide ion is oxidized by hydroxide ion.
d.
b.
MnO4¯ is oxidized by iodide ion.
e.
c.
The oxidation number of manganese changes
from +7 to +2.
The oxidation number of manganese remains
the same.
The oxidation number of iodine changes from
-1 to 0.
____ 189. Which of the following is true at the triple point of a pure substance?
a.
b.
c.
The vapor pressure of the solid phase always d.
equal the vapor pressure of the liquid phase.
The temperature is always 0.01 K lower that
e.
the normal melting point.
The liquid and gas phases of the substance
always have the same density and are therefore
indistinguishable.
the solid phase always melts if the pressure
increases at constant temperature.
The liquid phase always vaporizes if the
pressure increases at constant temperature.
____ 190. __ Cr2O72¯ + __ e¯ + __ H+ ---> __ Cr3+ + __ H2O(l) When the equation for the half reaction above is balanced with
the lowest whole-number coefficients, the coefficient for H2O is
a. 2
d. 7
b. 4
e. 14
c. 6
____ 191. How many grams of calcium nitrate, Ca(NO3)2, contains 24 grams of oxygen atoms?
a. 164 grams
d. 50. grams
b. 96 grams
e. 41 grams
c. 62 grams
____ 192. The mass of element Q found in 1.00 mole of each of four different compounds is 38.0 grams, 57.0 grams, 76.0
grams, and 114 grams, respectively. A possible atomic weight of Q is
a. 12.7
d. 38.0
b. 19.0
e. 57.0
c. 27.5
____ 193. The simplest formula for an oxide of nitrogen that is 36.8 percent nitrogen by weight is
a. N2O
d. N2O3
b. NO
e. N2O5
c. NO2
____ 194. How many milliliters of 11.6-molar HCl must be diluted to obtain 1.0 liter of 3.0-molar HCl?
a. 3.9 mL
d. 1,000 mL
b. 35 mL
e. 3,900 mL
c. 260 mL
____ 195. I. Difference in temperature between freezing point of solvent and freezing point of solvent and freezing point of
solution
II. Molal freezing point depression constant, Kf, for solvent
In addition to the information above, which of the following gives the minimum data required to determine the
molecular mass of a nonionic substance by the freezing point depression technique?
a.
b.
No further information is necessary.
Mass of solute
c.
Mass of solute and mass of solvent
d.
e.
Mass of solute and volume of solvent
Mass of solute, mass of solvent, and vapor
pressure of solvent
____ 196. Which of the following is probably true for a solid solute with a highly endothermic heat of solution when dissolved
in water?
a. The solid has a low lattice energy.
d. The solid is more soluble at higher
temperatures.
b. As the solute dissolves, the temperature of the e. the solid has a high energy of hydration.
solution increases.
c. The resulting solution is ideal.
____ 197. Hydrogen gas is collected over water at 24 °C. The total pressure of the sample is 755 millimeters of mercury. At 24
°C, the vapor pressure of water is 22 millimeters of mercury. What is the partial pressure of the hydrogen gas?
a. 22 mm Hg
d. 760 mm Hg
b. 733 mm Hg
e. 777 mm Hg
c. 755 mm Hg
____ 198. The structural isomers C2H5OH and CH3OCH3 would be expected to have the same values for which of the
following? (Assume ideal behavior.)
a. Gaseous densities at the same temperature and d. Melting points
pressure
b. Vapor pressures at the same temperature
e. Heats of vaporization
c. Boiling points
____ 199. A 2.00-liter sample of nitrogen gas at 27 °C and 600. millimeters of mercury is heated until it occupies a volume of
5.00 liters. If the pressure remains unchanged, the final temperature of the gas is
a. 68 °C
d. 677 °C
b. 120 °C
e. 950. °C
c. 477 °C
____ 200. Which of the following conclusions can be drawn from J. J. Thomson's cathode ray experiments?
a. Atoms contain electrons.
d. Atoms have a positively charged nucleus
surrounded by an electron cloud.
b. Practically all the mass of an atom is contained e. No two electrons in one atom can have the
in its nucleus.
same four quantum numbers.
c. Atoms contain protons, neutrons, and
electrons.
____ 201. Appropriate laboratory procedures include which of the following?
I. Rinsing a buret with distilled water just before filling it with the titrant for the first titration
II. Lubricating glass tubing before inserting it into a stopper
III. For accurate results, waiting until warm or hot objects have reached room temperature before weighing them
a. II only
d. II and III only
b. I and II only
e. I, II, and III
c. I and III only
____ 202. __ CH3CH2COOH(l) + __ O2(g) ---> __ CO2(g) + __ H2O(l)
How many moles of O2 are required to oxidize 1 mole of CH3CH2COOH according to the reaction represented
above?
a.
b.
c.
2 moles
5/2 moles
3 moles
d.
e.
7/2 moles
9/2 moles
____ 203. When a hydrate of Na2CO3 is heated until all the water is removed, it loses 54.3 percent of its mass. The formula of
the hydrate is
a. Na2CO3 . 10 H2O
d. Na2CO3 . 3 H2O
.
b. Na2CO3 7 H2O
e. Na2CO3 . H2O
c. Na2CO3 . 5 H2O
____ 204. 2 K + 2 H2O ---> 2 K+ + 2 OH¯ + H2
When 0.400 mole of potassium reacts with excess water at standard temperature and pressure as shown in the
equation above, the volume of hydrogen gas produced is
a.
b.
c.
1.12 liters
2.24 liters
3.36 liters
d.
e.
4.48 liters
6.72 liters
____ 205. Which of the following reactions has the largest positive value of S per mole of Cl2
a. H2(g) + Cl2(g) ---> 2 HCl(g)
d. 2 NH4Cl(s) ---> N2(g) + 4 H2(g) + Cl2(g)
b. Cl2(g) + 1/2 O2(g) ---> Cl2O(g)
e. Cl2(g) ---> 2 Cl(g)
c. Mg(s) + Cl2(g) ---> MgCl2(s)
____ 206. The SbCl5 molecule has trigonal bipyramid structure. Therefore, the hybridization of Sb orbitals should be
a. sp2
d. dsp3
b. sp3
e. d2sp3
2
c. dsp
____ 207. The metal calcium reacts with molecular hydrogen to form a compound. All of the following statements concerning
this compound are true EXCEPT:
a. Its formula is CaH2.
d. When added to water, it reacts to produce H2
gas.
b. It is ionic.
e. When added to water, it forms an acidic
solution.
c. It is solid at room temperature
____ 208. A measured mass of an unreactive metal was dropped into a small graduated cylinder half filled with water. The
following measurements were made.
Mass of metal = 19.611 gramsVolume of water before addition of metal = 12.4 millilitersVolume of water after
addition of metal = 14.9 milliliters
The density of the metal should be reported as
a.
b.
c.
7.8444 grams per mL
7.844 grams per mL
7.84 grams per mL
d.
e.
7.8 grams per mL
8 grams per mL
____ 209. CCl4, CO2, PCl3, PCl5, SF6 Which of the following does not describe any of the molecules above?
a. Linear
d. Tetrahedral
b. Octahedral
e. Trigonal pyramidal
c. Square planar
____ 210. Which of the following is a graph that describes the pathway of reaction that is endothermic and has high activation
energy?
a. A
d. D
b. B
e. E
c. C
____ 211. The normal boiling point of the substanced represented by the phase diagram above is
a.
b.
c.
-15 °C
-10 °C
140 °C
d.
e.
greater than 140 °C
not determinable from the diagram
____ 212. The phase diagram above provides sufficient information for determining the
a.
b.
c.
entropy change on vaporization
conditions necessary for sublimation
deviations from ideal gas behavior of the gas
phase
d.
e.
latent heat of vaporization
latent heat of fusion
____ 213. For the substance represented in the diagram, which of the phases is most dense and which is least dense at - 15 °C.
a. Most: Solid, Least: Gas
d. Most: Liquid, Least: Gas
b. Most: Solid, Least: Liquid
e. The diagram gives no information about
densities.
c. Most: Liquid, Least: Solid
____ 214. The test for the presence of Ag+ in an unknown solution involves the treatment of the solver-ammonia complex with
dilute hydrochloric acid. The appearance of a white precipitate at this point indicates the presence of silver ion in the
original sample. The net ionic equation that represents this test is
a. Ag(NH4)4+ + 4 H+ <===> Ag(s) + 4 NH4+
d. Ag(NH3)4+ + Cl¯ <===> Ag(NH3)2Cl(s)
+
+
b. Ag(NH4)4 + Cl¯ <===> AgCl(s) + 4 NH4
e. Ag(NH3)4+ + 2 H+ + Cl#175; <===> AgCl(s) +
2 NH4+
c. Ag(NH3)4+ + 4 HCl <===> AgCl(s) + 4 NH4+
+ 3 Cl#175;
____ 215. rate = k[X]
For the reaction whose rate law is given above, a plot of which of the following is a straight line?
a.
b.
c.
[X] versus time
log [X] versus time
1/[X] versus time
d.
e.
[X] versus 1/time
log [X] versus 1/time
____ 216.
(CH3)3CCl(aq) + OH¯ ---> (CH3)3COH(aq) + Cl¯
For the reaction represented above, the experimental rate law is given as follows.
Rate = k [(CH3)3CCl]
If some solid sodium solid hydroxide is added to a solution that is 0.010-molar in (CH3)3CCl and 0.10-molar in
NaOH, which of the following is true? (Assume the temperature and volume remain constant.)
a.
Both the reaction rate and k increase.
d.
b.
Both the reaction rate and k decrease.
e.
c.
Both the reaction rate and k remain the same.
The reaction rate increases but k remains the
same.
The reaction rate decreases but k remains the
same.
____ 217. Which of the following compounds is ionic and contains both sigma and pi covalent bonds?
a. Fe(OH)3
d. NO2
b. HClO
e. NaCN
c. H2S
____ 218. When acidified K2Cr2O7 solution is added to Na2S solution, green Cr3+ ions and free S are formed. When acidified
K2Cr2O7 solution is added to NaCl, no change occurs. Of the substances involved in these reactions, which is the
best reducing agent?
a. K2Cr2O7
d. S
b. Na2S
e. NaCl
c. Cr3+
____ 219. As the temperature is raised from 20 ° C to 40 ° C, the average kinetic energy of neon atoms changes by a factor of
a. 1/2
d. 2
b. [square root of](313/293)
e. 4
c. 313/293
____ 220. Which of the following characteristics is common to elemental sulfur, chlorine, nitrogen, and carbon?
a. They are gaseous elements at room
d. They form ionic oxides.
temperature.
b. They have oxides that are acid anhydrides.
e. They react readily with hydrogen at room
temperature.
c. They have perceptible color at room
temperature.
____ 221. A solution is known to contain an inorganic salt of one of the following elements. The solution is colorless. The
solution contains a salt of
a. Cu
d. Ni
b. Mn
e. Zn
c. Fe
____ 222. BrO3¯ + 5 Br¯ + 6 H+ <===> 3 Br2 + 3 H2O
If 25.0 milliliters of 0.200-molar BrO3¯ is mixed with 30.0 milliliters of 0.450-molar Br¯ solution that contains a
large excess of H+, the amount of Br2 formed, according to the equation above, is
a.
b.
c.
5.00 x 10¯3 mole
8.10 x 10¯3 mole
1.35 x 10¯2 mole
d.
e.
1.50 x 10¯2 mole
1.62 x 10¯2 mole
____ 223. The specific rate constant k for radioactive element X is 0.023 min -1. What weight of X was originally present in a
sample if 40. grams is left after 60. minutes?
a. 10. grams
d. 120 grams
b. 20. grams
e. 160 grams
c. 80. grams
____ 224. A white solid is observed to be insoluble in water, insoluble in excess ammonia solution, and soluble in dilute HCl.
Which of the following compounds could the solid be?
a. CaCO3
d. AgCl
b. BaSO4
e. Zn(OH)2
c. Pb(NO3)2
____ 225. H2O(s) ---> H2O(l) When ice melts at its normal melting point, 273.16 K and 1 atmosphere, which of the following
is true for the process shown above?
a. H < 0, S > 0, V > 0
d. H > 0, S > 0, V > 0
b. H < 0, S < 0, V > 0
e. H > 0, S > 0, V < 0
c. H > 0, S < 0, V < 0
____ 226. A solution of toluene (molecular weight 92.1) in benzene (molecular weight 78.1) is prepared. The mole fraction of
toluene in the solution is 0.100. What is the molality of the solution?
a. 0.100 m
d. 1.28 m
b. 0.703 m
e. 1.42 m
c. 0.921 m
____ 227. How many moles of solid Ba(NO3)2 should be added to 300. milliliters of 0.20-molar Fe(NO3)3 to increase the
concentration of the NO3¯ ion to 1.0-molar? (Assume that the volume of the solution remains constant.)
a.
b.
c.
0.060 mole
0.12 mole
0.24 mole
d.
e.
0.30 mole
0.40 mole
____ 228. Adding water to some chemicals can be dangerous because large amounts of heat are liberated. Which of the
following does NOT liberate heat when water is added to it?
a. KNO3
d. H2SO4
b. NaOH
e. Na
c. CaO
____ 229. Can be used to predict that a gaseous carbon atom in its ground state is paramagnetic
a. Heisenberg uncertainty principle
d. Shielding effect
b. Pauli exclusion principle
e. Wave nature of matter
c. Hund's rule (principle of maximum
multiplicity)
____ 230. Explains the experimental phenomenon of electron diffraction
a. Heisenberg uncertainty principle
d. Shielding effect
b. Pauli exclusion principle
e. Wave nature of matter
c. Hund's rule (principle of maximum
multiplicity)
____ 231. Indicates that an atomic orbital can hold no more than two electrons
a.
b.
c.
Heisenberg uncertainty principle
Pauli exclusion principle
Hund's rule (principle of maximum
multiplicity)
d.
e.
Shielding effect
Wave nature of matter
____ 232. Predicts that it is impossible to determine simultaneously the exact position and the exact velocity of an electron
a.
b.
c.
Heisenberg uncertainty principle
Pauli exclusion principle
Hund's rule (principle of maximum
multiplicity)
d.
e.
Shielding effect
Wave nature of matter
____ 233. If the temperature increases from 10° C to 60° C at a constant pressure of 0.4 atmosphcre, which of the processes
occurs?
a. Sublimation
d. Fusion
b. Condensation
e. Freezing
c. Solvation
____ 234. If the temperature decreases from 110° C to 40° C at a constant pressure of 1.1 atmospheres, which of the processes
occurs?
a. Sublimation
d. Fusion
b. Condensation
e. Freezing
c. Solvation
____ 235. If the pressure increases from 0.5 to 1.5 atmospheres at a constant temperature of 50° C, which of the processes
occurs?
a. Sublimation
d. Fusion
b. Condensation
e. Freezing
c. Solvation
____ 236. Has the largest bond-dissociation energy
a. Li2
b. B2
c. N2
d.
e.
O2
F2
____ 237. Has a bond order of 2
a. Li2
b. B2
c. N2
d.
e.
O2
F2
d.
e.
O2
F2
d.
e.
As
Na
____ 238. Contains 1 sigma (s) and 2 pi (p) bonds
a.
b.
c.
Li2
B2
N2
____ 239. Utilized as a coating to protect Fe from corrosion
a. Pb
b. Ca
c. Zn
____ 240. Is added to silicon to enhance its properties as a semiconductor
a.
b.
c.
Pb
Ca
Zn
____ 241. Utilized as a shield from sources of radiation
a. Pb
b. Ca
c. Zn
d.
e.
As
Na
d.
e.
As
Na
____ 242. In a molecule in which the central atom exhibits sp3d2 hybrid orbitals, the electron pairs are directed toward the
corners of
a. a tetrahedron
d. a square
b. a square-based pyramid
e. an octahedron
c. a trigonal bipyramid
____ 243. Relatively slow rates of chemical reaction are associated with which of the following?
a. The presence of a catalyst
d. Strong bonds in reactant molecules
b. High temperature
e. Low activation energy
c. High concentration of reactants
____ 244. 2 H2O + 4 MnO4¯ + 3 ClO2¯ ---> 4 MnO2 + 3 ClO4¯ + 4 OH¯
Which species acts as an oxidizing agent in the reaction represented above?
a.
b.
c.
H2O
ClO4¯
ClO2¯
d.
e.
MnO2
MnO4¯
____ 245. In which of the following compounds is the mass ratio of chromium to oxygen closest to 1.62 to 1.00 ?
a.
b.
c.
CrO3
CrO2
CrO
d.
e.
Cr2O
Cr2O3
____ 246. . . . Ag+ + . . . AsH3(g) + . . . OH¯ ---> . . . Ag(s) + . . . H3AsO3(aq) + . . . H2O
When the equation above is balanced with lowest whole-number coefficients, the coefficient for OH¯ is
a.
b.
c.
2
4
5
d.
e.
6
7
____ 247. Step 1: Ce4+ + Mn2+ ---> Ce3+ + Mn3+
Step 2: Ce4+ + Mn3+ ---> Ce3+ + Mn4+
Step 3: Mn4+ + Tl+ ---> Tl3+ + Mn2+
The proposed steps for a catalyzed reaction between Ce 4+ and Tl+ are represented above. The products of the overall
catalyzed reaction are
a.
b.
c.
Ce4+ and Tl+
Ce3+ and Tl3+
Ce3+ and Mn3+
d.
e.
Ce3+ and Mn4+
Tl3+ and Mn2+
____ 248. A sample of 0.0100 mole of oxygen gas is confined at 37° C and 0.216 atmosphere. What would be the pressure of
this sample at 15° C and the same volume?
a. 0.0876 atm
d. 0.233 atm
b. 0.175 atm
e. 0.533 atm
c. 0.201 atm
H° = - 286 kJ
H° = - 414 kJ
H° = - 425 kJ
____
249.
H2(g) + (1/2) O2(g) ---> H2O(l)
2 Na(s) + (1/2) O2(g) ---> Na2O(s)
Na(s) + (1/2) O2(g) + (1/2) H2(g) ---> NaOH(s)
Based on the information above, what is the standard enthalpy change for the following reaction?
Na2O(s) + H2O(l) ---> 2 NaOH(s)
a.
b.
c.
-1,125 kJ
-978 kJ
-722 kJ
d.
e.
-150 kJ
+275 kJ
____ 250. Which of the following actions would be likely to change the boiling point of a sample of a pure liquid in an open
container?
I. Placing it in a smaller container
II. Increasing the number of moles of the liquid in the container
III. Moving the container and liquid to a higher altitude
a.
b.
c.
I only
II only
III only
d.
e.
II and III only
I, II, and III
____ 251. Which of the following sets of quantum numbers (n, l, ml, ms) best describes the valence electron of highest energy
in a ground-state gallium atom (atomic number 31) ?
a. 4, 0, 0, 1/2
d. 4, 1, 2, 1/2
b. 4, 0, 1, 1/2
e. 4, 2, 0, 1/2
c. 4, 1, 1, 1/2
____ 252. Given that a solution is 5 percent sucrose by mass, what additional information is necessary to calculate the molarity
of the solution?
I. The density of water
II. The density of the solution
III. The molar mass of sucrose
a.
b.
c.
I only
II only
III only
d.
e.
I and III
II and III
____ 253. When an aqueous solution of NaOH is added to an aqueous solution of potassium dichromate, K 2Cr2O7 the
dichromate ion is converted to
a. CrO42¯
d. Cr2O3(s)
b. CrO2
e. Cr(OH)3(s)
c. Cr3+
____ 254. CH3CH2OH boils at 78 °C and CH3OCH3 boils at - 24 °C, although both compounds have the same composition.
This difference in boiling points may be attributed to a difference in
a. molecular mass
d. hydrogen bonding
b. density
e. heat of combustion
c. specific heat
____ 255. A hydrocarbon gas with an empirical formula CH2 has a density of 1.88 grams per liter at 0 °C and 1.00 atmosphere.
A possible formula for the hydrocarbon is
a. CH2
d. C4H8
b. C2H4
e. C5H10
c. C3H6
____ 256. X = CH3-CH2-CH2-CH2-CH3
Y = CH3-CH2-CH2-CH2-OH
Z = HO-CH2-CH2-CH2-OH
Based on concepts of polarity and hydrogen bonding, which of the following sequences correctly lists the
compounds above in the order of their increasing solubility in water?
a.
b.
c.
Z<Y<X
Y<Z<X
Y<X<Z
d.
e.
X<Z<Y
X<Y<Z
____ 257. A sample of 3.30 grams of an ideal gas at 150.0 °C and 1.25 atmospheres pressure has a volume of 2.00 liters. What
is the molar mass of the gas? The gas constant, R, is 0.0821 L atm mol¯ 1 K¯1).
a. 0.0218 gram/mole
d. 45.8 grams/mole
b. 16.2 grams/mole
e. 71.6 grams/mole
c. 37.0 grams/mole
____ 258. Concentrations of colored substances are commonly measured by means of a spectrophotometer. Which of the
following would ensure that correct values are obtained for the measured absorbance?
I. There must be enough sample in the tube to cover the entire light path.
II. The instrument must be periodically reset using a standard.
III. The solution must be saturated.
a.
b.
c.
I only
II only
I and II only
d.
e.
II and III only
I, II, and III
____ 259. Samples of F2 gas and Xe gas are mixed in a container of fixed volume. The initial partial pressure of the F 2 gas is
8.0 atmospheres and that of the Xe gas is 1.7 atmospheres. When all of the Xe gas reacted, forming a solid
compound, the pressure of the unreacted F2 gas was 4.6 atmospheres. The temperature remained constant. What is
the formula of the compound?
a. XeF
d. XeF6
b. XeF3
e. XeF8
c. XeF4
____ 260. A strip of metallic scandium, Sc, is placed in a beaker containing concentrated nitric acid. A brown gas rapidly
forms, the scandium disappears, and the resulting liquid is brown-yellow but becomes colorless when warmed.
These observations best support which of the following statements?
a. Nitric acid is a strong acid.
d. Scandium reacts with nitric acid to form a
brown gas.
b. In solution scandium nitrate is yellow and
e. Scandium and nitric acid react in mole propor
scandium chloride is color less.
tions of 1 to 3.
c. Nitric acid reacts with metals to form
hydrogen.
____ 261. Mass of an empty container = 3.0 grams
Mass of the container plus the solid sample = 25.0 grams
Volume of the solid sample = 11.0 cubic centimeters
The data above were gathered in order to determinethe density of an unknown solid. The density of the sample
should be reported as
a.
b.
c.
0.5 g/cm3
0.50 g/cm3
2.0 g/cm3
d.
e.
2.00 g/cm3
2.27 g/cm3
____ 262. Which of the following solutions has the lowest freezing point?
a. 0.20 m C6H12O6, glucose
d. 0.20 m KMnO4
b. 0.20 m NH4Br
e. 0.20 m MgCl2
c. 0.20 m ZnSO4
____ 263. A sample of an ideal gas is cooled from 50.0 °C to 25.0 °C in a sealed container of constant volume. Which of the
following values for the gas will decrease?
I. The average molecular mass of the gas
II. The average distance between the molecules
III. The average speed of the molecules
a.
b.
c.
I only
II only
III only
d.
e.
I and III
II and III
____ 264. Which of the following solids dissolves in water to form a colorless solution?
a. CrCl3
d. CuCl2
b. FeCl3
e. ZnCl2
c. CoCl2
____ 265. Which of the following has the lowest conductivity?
a. 0.1 M CuS04
d.
b. 0.1 M KOH
e.
c. 0.1 M BaCl2
0.1 M HF
0.1 M HNO3
____ 266. The isomerization of cyclopropane to propylene is a first-order process with a half-life of 19 minutes at 500 °C. The
time it takes for the partial pressure of cyclopropane to decrease from 1.0 atmosphere to 0.125 atmosphere at 500 °C
is closest to
a. 38 minutes
d. 152 minutes
b. 57 minutes
e. 190 minutes
c. 76 minutes
____ 267. Which of the following acids can be oxidized to form a stronger acid?
a. H3PO4
d. H3BO3
b. HNO3
e. H2SO3
c. H2CO3
____ 268. When dilute nitric acid was added to a solution of one of the following chemicals, a gas was evolved, This gas
turned a drop of limewater, Ca(OH)2, cloudy, due to the formation of a white precipitate. The chemical was
a. household ammonia, NH3
d. epsom salts, MgSO4 . 7H2O
b. (B) baking soda, NaHCO3
e. bleach, 5% NaOCl
c. table salt, NaCl