Download First Semester Final Review

First Semester Final Review
1. The energy required to convert a ground-state atom in the gas phase to a gaseous positive ion
a. Activation energy b. Free energy c. Ionization energy d. Kinetic energy e. Lattice energy
2. The energy change that occurs in the conversion of an ionic solid to widely separated gaseous ions
a. Activation energy b. Free energy c. Ionization energy d. Kinetic energy e. Lattice energy
3. Represents an atom that is chemically unreactive
a. 1s ___ 2s b. 1s  2s c. 1s  2s  2p   _ d. 1s  2s  2p   e. [Ar] 4s  3d
    
4. Represents an atom in an excited state
a. 1s ___ 2s 
b. 1s  2s 
c. 1s  2s  2p   _
d.
e.
1s  2s  2p   
[Ar] 4s  3d     
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   __
d.
e.
1s  2s  2p   
[Ar] 4s  3d     
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
d. Strong multiple covalent bonds (including p-bonds) with weak intermolecular forces e. Macromolecules
held together with strong polar bonds
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
d. Strong multiple covalent bonds (including p-bonds) with weak intermolecular forces e. Macromolecules
held together with strong polar bonds
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
d. Strong multiple covalent bonds (including p-bonds) with weak intermolecular forces e. Macromolecules
held together with strong polar bonds
10. Methane, CH4(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
d. Strong multiple covalent bonds (including p-bonds) with weak intermolecular forces e. Macromolecules
held together with strong polar bonds
11. Is a gas in its standard state at 298 K
a. Lithium b. Nickel c. Bromine d. Uranium e. Flourine
12. Reacts with water to form a strong base
a. Lithium b. Nickel c. Bromine d. Uranium e. Flourine
13. What mass of Au is produced when 0.0500 mol of Au2S3 is reduced completely with excess H2?
a. 9.85 g b. 19.7 g c. 24.5 g d. 39.4 g e. 48.9 g
14. When a solution of sodium chloride is vaporized in a flame, the color of the flame is
a. blue b. yellow c. green d. violet e. white
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 increases with increasing temperature. b. The difference in
temperature between the air inside and outside the balloon produces convection currents. c. The cooler
air outside the balloon pushes in on the walls of the balloon. d. The rate of diffusion of cooler air is less
than that of warmer air. e. The air density inside the balloon is less than that of the surrounding air.
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 b. Sprinkle the affected area with powdered Na2SO4(s)
c. Flush the affected area with water and then with a dilute NaOH solution d. Flush the affected area with
water and then with a dilute NaHCO3 solution e. Flush the affected area with water and then with a dilute
vinegar 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 b. point R only c. all points on the curve between Q and S d. all points on the curve
between R and T e. no point on the curve
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 b. 7 c. 12 d. 14 e. 28
19. 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 b. I and II only c. I and III only d. II and III only e. I, II, and III
CH 3
O
||
 C  CH 2  CH 3
20. The organic compound represented above is an example of
a. an organic acid b. an alcohol c. an ether d. an aldehyde e. a ketone
21. 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 change. b. The oxidation number of H changes from -1 to +1.
c. The oxidation number of F changes from +1 to -1. d. The oxidation number of Se changes from -2 to
+6. e. It is a disproportionation reaction for F.
22. 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 unchanged. b. The molarity of the solution remains unchanged.
c. The molality of the solution remains unchanged. d. The mole fraction of the solute decreases. e. The
mole fraction of solute increases.
23. 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 b. III only c. I and II only d. II and III only e. I, II, and III
24. 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. 0.10 mol b. 0.20 mol c. 0.30 mol d. 0.40 mol e. 0.60 mol
25. 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 b. Mg c. Al d. Si e. P
26. 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 b. B c. C d. D e. E
27. Of the following molecules, which has the largest dipole moment?
a. CO b. CO2 c. O2 d. HF e. F2
28. ___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 b. 2 c. 3 d. 4 e. 6
29. 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. b. The pressure of the gas decreases. c. The average speed of the
gas molecules remains the same. d. The total number of gas molecules remains the same. e. The
average distance between the gas molecules increases.
30. Which of the following occurs when excess concentrated NH 3 (aq) is mixed thoroughly with 0.1 M
Cu(NO3)2 (aq) ?
a. A dark red precipitate forms and settles out. b. Separate layers of the immiscible liquids form with a
blue layer on top. c. The color of the solution turns from light blue to dark blue. d. Bubbles of ammonia
gas form. e. The pH of the solution decreases.
31. 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 b. HfCl2 c. HfCl3 d. HfCl4 e. Hf2Cl3
32. 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 b. Filtration c. Titration d. Electrolysis e. Evaporation to dryness
33. In the periodic table, as the atomic number increases from 11 to 17, what happens to the atomic radius?
a. It remains constant. b. It increases only. c. It increases, then decreases. d. It decreases only. e. It
decreases, then increases.
34. 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 negative charges. b. Electrons in atoms are arranged in
shells. c. Neutrons are at the center of an atom. d. Neutrons and protons in atoms have nearly equal
mass. e. The positive charge of an atom is concentrated in a small region.
35. 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. A b. B c. C d. D e. E
36. 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 b. 0.40 atm c. 1.0 atm d. 1.2 atm e. 1.4 atm
37. 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. b. 10. c. 8.0 d. 5.0 e. 2.5
38. 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) b. Zn2+(aq) c. CrO42-(aq) d. SO42-(aq) e. OH-(aq)
39. 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 due to the higher atmospheric pressure. b. equilibrium
water vapor pressure is lower due to the higher atmospheric pressure. c. equilibrium water vapor
pressure equals the atmospheric pressure at a lower temperature d. water molecules have a higher
average kinetic energy due to the lower atmospheric pressure e. water contains a greater concentration
of dissolved gases
40. 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 b. 0.19 M c. 0.28 M d. 0.40 M e. 0.55 M
41. 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 b. 1 atm c. 0.5 atm d. 0.1 atm e. 0.03 atm
42. 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 b. -1,279 kJ c. -1,323 kJ d. -1,367 kJ e. -1,411 kJ
43. 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 b. PHe<PAr<PNe c. PNe<PAr<PHe d. PAr<PHe< PNe e. PAr=PHe= PNe
44. Which of the following compounds is NOT appreciably soluble in water but is soluble in dilute hydrochloric
acid?
a. Mg(OH)2(s) b. (NH4)2CO3(s) c. CuSO4(s) d. (NH4)2SO4(s) e. Sr(NO3)2(s)
45. In which of the following processes are covalent bonds broken?
a. I2(s)  I2(g) b. CO2(s)  CO2(g) c. NaCl(s)  NaCl(l) d. C(diamond)  C(g) e. Fe(s)  Fe(l)
46. 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 b. 0.012 M c. 0.025 M d. 0.075 M e. 0.10 M
47. 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+] b. [PO43-] < [Na+] < [NO3-] c. [NO3-] < [PO43-] < [Na+] d. [Na+] < [NO3-] < [PO43-]
e. [Na+] < [PO43-] < [NO3-]
48. 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 b. Adding dilute HCl(aq) solution c. Adding dilute NaOH(aq) solution
d. Adding dilute NH3(aq) solution e. Adding dilute HNO3(aq) solution
49. 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 hydrate sample to spatter out of the crucible. b. The
dehydrated sample absorbed moisture after heating. c. The amount of the hydrate sample used was too
small. d. The crucible was not heated to constant mass before use. e. Excess heating caused the
dehydrated sample to decompose.
50. 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 b. 60.0 mL c. 100. mL d. 110. mL e. 120. mL
51. Which of the following gases deviates most from ideal behavior?
a. SO2 b. Ne c. CH4 d. N2 e. H2
52. 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) b. CH3CH2CH2OH(l) and H2O(l) c. CH3CH2CH2OH(l) and C8H18(l) d. C6H14(l) and
C8H18(l) e. H2SO4(l) and H2O(l)
53. The atom that contains exactly two unpaired electrons
a. S b. Ca c. Ga d. Sb e. Br
54. The atom that contains only one electron in the highest occupied energy sublevel
a. S b. Ca c. Ga d. Sb e. Br
55. The molecule with only one double bond
a. CO2 b. H2O c. CH4 d. C2H4 e. PH3
56. The molecule with the largest dipole moment
a. CO2 b. H2O c. CH4 d. C2H4 e. PH3
57. The molecule that has trigonal planar pyramidal geometry
a. CO2 b. H2O c. CH4 d. C2H4 e. PH3
58. Is purple in aqueous solution
a. PbSO4 b. CuO c. KMnO4 d. KCl e. FeCl3
59. Is white and very soluble in water
a. PbSO4 b. CuO c. KMnO4 d. KCl e. FeCl3
60. Has an average atomic or molecular speed closest to that of N2 molecules at 0o and 1 atm
a. Ne b. Xe c. O2 d. CO e. NO
61. Has the greatest density
a. Ne b. Xe c. O2 d. CO e. NO
62. Has the greatest rate of effusion through a pinhole
a. Ne b. Xe c. O2 d. CO e. NO
63. 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-(aq) + 3 H2O(l) d. Ca2+(aq) + SO42-(aq)  CaSO4(s) e. PtCl4(s) +
2 Cl-(aq)  PtCl6264. A reaction in which the same reactant undergoes both oxidation and reduction
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)  PtCl6265. 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-
66. The substance is at its normal freezing point at time
a. t1 b. t2 c. t3 d. t4 e. t5
67. Which of the following best describes what happens to the substance between t4 and t5?
a. The molecules are leaving the liquid phase. b. The solid and liquid phases coexist in equilibrium.
c. The vapor pressure of the substance is decreasing. d. The average intermolecular distance is
decreasing. e. The temperature of the substance is increasing.
68. In which of the following groups are the three species isoelectronic; i.e., have the same number of
electrons?
a. S2-, K+, Ca2+ b. Sc, Ti, V2+ c. O2-, S2-, Cl- d. Mg2+, Ca2+, Sr2+ e. Cs, Ba2+, La3+
69. 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 about 70 oC. b. It first melts to a liquid and then boils at about
30oC. c. It melts to a liquid at a temperature of about 20oC and remains a liquid until the temperature is
greater than 100oC. d. It sublimes to a vapor at an equilibrium temperature of about 20oC. e. It remains
a solid until the temperature is greater than 100oC.
70. In which of the following species does sulfur have the same oxidation number as it does in H2SO4?
a. H2SO3 b. S2O32- c. S2- d. S8 e. SO2Cl2
71. 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 SO 2(g) in the flask?
a. 800 mm Hg b. 600 mm Hg c. 250 mm Hg d. 200 mm Hg e. 160 mm Hg
72. 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 b. II only c. III only d. I and II only e. I, II, and III
73. 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 b. KTe2O c. K2TeO3 d. K2TeO6 e. K4TeO6
74. 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 b. -147 kJ c. -19 kJ d. +19 kJ e. +773 kJ
75. 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 b. 6.2 g c. 34 g d. 85 g e. 140 g
76. Of the following compounds, which is the most ionic?
a. SiCl4 b. BrCl c. PCl3 d. Cl2O e. CaCl2
77. The best explanation for the fact that diamond is extremely hard is that diamond crystals
a. are made up of atoms that are intrinsically hard because of their electronic structures 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 d. are formed under extreme
conditions of temperature and pressure e. contain orbitals or bands of delocalized electrons that belong
not to single atoms but to each crystal as a whole
78. 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 b. 22.4 L c.
X 22.4 L d. 2 X 22.4 L e. 3 X 22.4 L
79. Which of the following oxides is a gas at 25oC and 1 atm?
a. Rb2O b. N2O c. Na2O2 d. SiO2 e. La2O3
80. 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 b. has a higher vapor pressure c. has the same vapor pressure d. has a
higher freezing point e. is more nearly ideal
81. 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) b. Ag+(aq) and Pb2+(aq) c. Ba2+(aq) and Ag+(aq) d. Ba2+(aq) and Hg22+(aq)
e. Ba2+(aq) and Pb2+(aq)
82. 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 NaCl is. b. Na2SO4 has a higher molar mass than NaCl has.
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. d. More moles of ions are present in a given volume of 0.10 M Na2SO4 than in the same
volume of 0.10 M NaCl. e. The degree of dissociation of Na2SO4 in solution is significantly greater than
that of NaCl.
83. 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 b. KNO3 c. K2CrO4 d. K2SO4 e. Ce2(SO4)3
84. 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 b. 44.8 mL c. 224 mL d. 448 mL e. 896 mL
85. Which of the following properties generally decreases across the periodic table from sodium to chlorine?
a. First ionization energy b. Atomic mass c. Electrongeativity d. Maximum value of oxidation number
e. Atomic radius
86. 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 b. 0.46 c. 0.54 d. 0.67 e. 0.75
87. 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 conditions than Ne. b. Na has a lower first ionization energy than
Ne. c. Na has a higher melting point than Ne. d. Na has a higher neutron-to-proton ratio than Ne. e. Na
has a fewer naturally occuring isotopes than Ne.
88. Sodium chloride is LEAST soluble in which of the following liquids?
a. H2O b. CCl4 c. HF d. CH3OH e. CH3COOH
89. ...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 b. 4 c. 6 d. 8 e. 14
90. Which of the following represents acceptable laboratory practice?
a. Placing a hot object on a balance pan b. Using distilled water for the final rinse of a buret before filling
it with standardized solution c. Adding a weighed quantity of solid acid to a titration flask wet with distilled
water d. Using 10 mL of standard strength phenolphthalein indicator solution for titration of 25 mL of acid
solution e. Diluting a solution in a volumetric flask to its final concentration with hot water
91. 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 b. II only c. III only d. I and II e. II and III
92. 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 b. 5 c. 7 d. 10 e. 22
93. 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 b. decreasing size of the central atom c. increasing electronegativity
of the central atom d. increasing the number of unshared pairs of electrons e. decreasing repulsion
between hydrogen atoms
94. 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. b. The hydrogen bonding increases. c. The dipole-dipole
forces increase. d. The chemical reactivity increases. e. The number of nearest neighbors increases.
95. 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 b. 18 g c. 36 g d. 72 g e. 144 g
96. All of the halogens in their elemental form at 25oC and 1 atm are
a. conductors of electricity b. diatomic molecules c. odorless d. colorless e. gases
97. 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 b. 0.0053 mol c. 0.0075 mol d. 0.013 mol e. 0.030 mol
98. Which of the following substances is LEAST soluble in water?
a. (NH4)2SO4 b. KMnO4 c. BaCO3 d. Zn(NO3)2 e. Na3PO4
99. A 2 L container will hold about 4 g of which of the following gases at 0oC and 1 atm?
a. SO2 b. N2 c. CO2 d. C4H8 e. NH3
100. 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 form. b. Hydrogen bonds between H2O molecules are
broken. c. Covalent bonds between H2O molecules are broken. d. Ionic bonds between H+ ions and
OH- ions are broken. e. Covalent bonds between H+ ions and H2O molecules become more effective.
101. 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 b. B c. C d. D e. E
102. 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 b. 0.30 M c. 0.40 M d. 0.60 M e. 1.2 M
103. Of the following pure substances, which has the highest melting point?
a. S8 b. I2 c. SiO2 d. SO2 e. C6H6
104. 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) b. Al3+(aq) c. Ba2+(aq) d. Na+(aq) e. CO32-(aq)
105. Forms monatomic ions with 2¯ charge in solutions
a. F b. S c. Mg d. Ar e. Mn
106. Forms a compound having the formula KXO4
a. F b. S c. Mg d. Ar e. Mn
107. Forms oxides that are common air pollutants and that yield acidic solution in water
a. F b. S c. Mg d. Ar e. Mn
108. Is a good oxidizing agent
a. Hydrofluoric acid b. Carbon dioxide c. Aluminum hydroxide d. Ammonia e. Hydrogen peroxide
109. Is used to etch glass chemically
a. Hydrofluoric acid b. Carbon dioxide c. Aluminum hydroxide d. Ammonia e. Hydrogen peroxide
110. Is used extensively for the production of fertilizers
a. Hydrofluoric acid b. Carbon dioxide c. Aluminum hydroxide d. Ammonia e. Hydrogen peroxide
111. Solid ethyl alcohol, C2H5OH
a. A network solid with covalent bonding b. A molecular solid with zero dipole moment c. A molecular
solid with hydrogen bonding d. An ionic solid e. A metallic solid
112. Silicon dioxide, SiO2
a. A network solid with covalent bonding b. A molecular solid with zero dipole moment c. A molecular
solid with hydrogen bonding d. An ionic solid e. A metallic solid
113. 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. Ba2+ e. Al3+
114. 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¯ b. Cr2O72¯ c. NH4+ d. Ba2+ e. Al3+
115. 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. CO32¯ b. Cr2O72¯ c. NH4+ d. Ba2+ e. Al3+
116. 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¯ b. Cr2O72¯ c. NH4+ d. Ba2+ e. Al3+
117.
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. HF gas is more ideal. b. HF is the strongest acid. c. HF molecules have a smaller dipole moment.
d. HF is much less soluble in water. e. HF molecules tend to form hydrogen bonds.
118. 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
119. .....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. 4 b. 6 c. 8 d. 9 e. 10
120. 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. The density of the gas b. The pressure of the gas c. The average velocity of the gas molecules
d. The number of molecules per cm 3 e. The potential energy of the molecules
121. 1s2 2s22p6 3s23p3
Atoms of an element, X, have the electronic configuration shown above. The compound most likely
formed with magnesium, Mg, is
a. MgX b. Mg2X c. MgX2 d. MgX3 e. Mg3X2
122. 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 b. 46.0 c. 88.0 d. 94.1 e. 138
123. 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. temperature at which the solid, liquid,
and vapor phases are all in equilibrium 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
124. A 0.1-molar solution of which of the following ions is orange?
a. Fe(H2O)42+ b. Cu(NH3)42+ c. Zn(OH)42¯ d. Zn(NH3)42+ e. Cr2O72¯
125. The net ionic equation for the reaction between silver carbonate and hydrochloric acid is
a. Ag2CO3(s) + 2 H+ + 2 Cl¯ ---> 2 AgCl(s) + H2O + CO(g) b. 2 Ag+ + CO32¯ + 2 H+ + 2 Cl¯ ---> 2 AgCl(s)
+ H2O + CO2(g) c. CO32¯ + 2 H+ ---> H2O + CO2(g) d. Ag+ + Cl¯ ---> AgCl(s) e. Ag2CO3(s) + 2 H+ --->
2Ag+ + H2CO3
126. ...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 b. 2:1 c. 3:1 d. 4:1 e. 5:1
127. 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+ b. Br¯ c. NO3¯ d. I¯ e. Co2+
128. The molality of the glucose in a 1.0-molar glucose solution can be obtained by using which of the
following?
a. Volume of the solution b. Temperature of the solution c. Solubility of glucose in water d. Degree of
dissociation of glucose e. Density of the solution
129. 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 b. It depends on the intermolecular forces of attraction between
molecules of X, Y, and Z. c. It depends on the relative molecular masses of X, Y, and Z. d. It depends
on the average distance traveled between molecular collisions. e. It can be calculated with knowledge
only of the volume of the container.
130. The geometry of the SO3 molecule is best described as
a. trigonal planar b. trigonal pyramidal c. square pyramidal d. bent
e. tetrahedral
131. Which of the following molecules has the shortest bond length?
a. N2 b. O2 c. Cl2 d. Br2 e. I2
132. Metallic copper is heated strongly with concentrated sulfuric acid. The products of this reaction are
a. CuSO4(s) and H2(g) only b. Cu2+, SO2(g), and H2O c. Cu2+, H2(g), and H2O d. CuSO4(s), H2(g), and
SO2(g) e. Cu2+, SO3(g), and H2O
133. The elements in which of the following have most nearly the same atomic radius?
a. Be, B, C, N b. Ne, Ar, Kr, Xe c. Mg, Ca, Sr, Ba d. C, P, Se, I e. Cr, Mn, Fe, Co
134. 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 b. 0.0625 mole c. 0.125 mole d. 0.250 mole e. 0.500 mole
135. The alkenes are compounds of carbon and hydrogen with the general formula C nH2n. If 0.561 gram of any
alkene is burned in excess oxygen, what number of moles of H 2O is formed?
a. 0.0400 mole b. 0.0600 mole c. 0.0800 mole d. 0.400 mole e. 0.800 mole
136. 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. -210.0 kJ/mole b. -107.5 kJ/mole c. -75.8 kJ/mole d. 75.8 kJ/mole e. 210.0 kJ/mole
137. 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 the same as the volume of the oxygen container. b. The
number of molecules in the hydrogen container is the same as the number of molecules in the oxygen
container. c. The density of the hydrogen sample is less than that of the oxygen sample. d. The
average kinetic energy of the hydrogen molecules is the same as the average kinetic energy of the
oxygen molecules. e. The average speed of the hydrogen molecules is the same as the average speed
of the oxygen molecules.
138. Pi bonding occurs in each of the following species EXCEPT
a. CO2 b. C2H4 c. CN¯ d. C6H6 e. CH4
139. 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. 0.015 mole b. 0.020 mole c. 0.030 mole d. 0.045 mole e. 0.090 mole
140. Which of the following statements is always true about the phase diagram of any one-component
system?
a. The slope of the curve representing equilibrium between the vapor and liquid phases is positive.
b. The slope of the curve representing equilibrium between the liquid and solid phases is negative.
c. The slope of the curve representing equilibrium between the liquid and solid phases is positive. d. the
temperature at the triple point is greater than the normal freezing point. e. The pressure at the triple point
is greater than 1 atmosphere.
141. At 20. °C, the vapor pressure of toluene is 22 millimeters of mercury and that of benzene is 75 millimeters
of mercury. An ideal solution, equimolar in toluene and benzene, is prepared. At 20. °C, what is the mole
fraction of benzene in the vapor in equilibrium with this solution?
a. 0.23 b. 0.29 c. 0.50 d. 0.77 e. 0.83
142. Which of the following represents the ground state electron configuration for the Mn 3+ ion? (Atomic
number Mn = 25)
a. 1s2 2s22p6 3s23p63d4 b. 1s2 2s22p6 3s23p63d5 4s2 c. 1s2 2s22p6 3s23p63d2 4s2 d. 1s2 2s22p6 3s23p63d8
4s2 e. 1s2 2s22p6 3s23p63d3 4s1
143. When 70. milliliter of 3.0-molar Na2CO3 is added to 30. milliliters of 1.0-molar NaHCO3 the resulting
concentration of Na+ is
a. 2.0 M b. 2.4 M c. 4.0 M d. 4.5 M e. 7.0 M
144. Which of the following has a zero dipole moment?
a. HCN b. NH3 c. SO2 d. NO2 e. PF5
145. 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 b. FeCrO4(s) and H2O c. Fe3+, CrO42¯, and H2O d. Fe3+, Cr3+, and H2O
e. Fe2(SO4)3(s), Cr3+ and H2O
146. 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. The burette was not rinsed with NaOH solution. b. The student misread a 5 for a 6 on the burette when
the first sample was titrated. c. A different amount of water was added to the first sample. d. The pipette
was not rinsed with the HCI solution. e. The student added too little indicator to the first sample.
147. 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 b. HNO2 + NaOH ---> Na+ + NO2¯ + H2O c. H+ + OH¯ --->H2O
d. HNO2 + H2O ---> NO2¯ + H3O+ e. HNO2 + OH¯ ---> NO2¯ + H2O
148. Which of the following species CANNOT function as an oxidizing agent?
a. Cr2O72¯ b. MnO4¯ c. NO3¯ d. S e. I¯
149. Ca, V, Co, Zn, As
Gaseous atoms of which of the elements above are paramagnetic?
a. Ca and As only b. Zn and As only c. Ca, V, and Co only d. V, Co, and As only e. V, Co, and Zn only
150. 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 kilograms of H2O b. 42.8 grams of KIO3 and add H2O until the final
homogeneous solution has a volume 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 d. 42.8 grams of KIO3 and add 2.00 liters of H2O
e. 21.4 grams fo KIO3 and add 2.00 liters of H2O
151. 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 b. 0.160 M c. 0.200 M d. 0.240 M e. 0.267 M
152. What is the mole fraction of ethanol, C2H5OH, in an aqueous solution in which the ethanol concentration
is 4.6 molal?
a. 0.0046 b. 0.076 c. 0.083 d. 0.20 e. 0.72
153. 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 b. 5, 1, 1, 1/2 c. 5, 1, 0, 1/2 d. 5, 0, 1, 1/2 e. 5, 0, 0, 1/2
154. 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 experiment b. Temperature of the water in the trough c. Vapor
pressure of the water d. Barometric pressure e. Volume of water displaced from the flask
155. 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 b. C2H2 c. C4H3 d. C4H6 e. C4H10
156. 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 b. mass c. velocity d. attractions e. shape
157. 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. 0.0010 M b. 0.0056 M c. 0.028 M d. 0.090 M e. 0.14 M
158. For which of the following molecules are resonance structures necessary to describe the bonding
satisfactorily?
a. H2S b. SO2 c. CO2 d. OF2 e. PF3
159. 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¯ ---> (NH4)2SO4 + Cu(OH)2 b. Cu2+ + 4 NH3 + 4 H2O --> Cu(OH)42¯ + 4
NH4+ c. Cu2+ + 2 NH3 + 2 H2O --> Cu(OH)2 + 2 NH4+ d. Cu2+ + 4 NH3 --> Cu(NH3)42+ e. Cu2+ + 2 NH3 +
H2O --> CuO + 2 NH4+
160. Which of the following aqueous solutions has the highest boiling point?
a. 0.10 M potassium sulfate, K2SO4 b. 0.10 M hydrochloric acid, HCl c. 0.10 M ammonium nitrate,
NH4NO3 d. 0.10 M magnesium sulfalte, MgSO4 e. 0.20 M sucrose, C12H22O11
161. 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 b. 11.2 liters c. 7.46 liters d. 5.60 liters e. 3.74 liters
162. What is the most electronegative element of the below?
a. O b. La c. Rb d. Mg e. N
163. Which element exhibits the greatest number of different oxidation states?
a. O b. La c. Rb d. Mg e. N
164. Which of the elements above has the smallest ionic radius for its most commonly found ion?
a. O b. La c. Rb d. Mg e. N
165. An impossible electronic configuration
a. 1s2 2s22p5 3s23p5 b. 1s2 2s22p6 3s23p6 c. 1s2 2s22p62d10 3s23p6 d. 1s2 2s22p6 3s23p63d5 e. 1s2 2s22p6
3s23p63d3 4s2
166. The ground-state configuration for the atoms of a transition element
a. 1s2 2s22p5 3s23p5 b. 1s2 2s22p6 3s23p6 c. 1s2 2s22p62d10 3s23p6 d. 1s2 2s22p6 3s23p63d5 e. 1s2 2s22p6
3s23p63d3 4s2
167. The ground-state configuration of a negative ion of a halogen
a. 1s2 2s22p5 3s23p5 b. 1s2 2s22p6 3s23p6 c. 1s2 2s22p62d103s23p6 d. 1s2 2s22p6 3s23p63d5 e. 1s2 2s22p6
3s23p63d34s2
168. The ground-state configuration of a common ion of an alkaline earth element
a. 1s2 2s22p5 3s23p5 b. 1s2 2s22p6 3s23p6 c. 1s2 2s22p62d10 3s23p6 d. 1s2 2s22p6 3s23p63d5 e. 1s2 2s22p6
3s23p63d3 4s2
169. Is used to explain why iodine molecules are held together in the solid state
a. hydrogen bonding b. hybridization c. ionic bonding d. resonance e. van der Waals forces (London
dispersion forces)
170. Is used to explain why the boiling point of HF is greater than the boiling point of HBr
a. hydrogen bonding b. hybridization c. ionic bonding d. resonance e. van der Waals forces (London
dispersion forces)
171. Is used to explain the fact that the four bonds in methane are equivalent
a. hydrogen bonding b. hybridization c. ionic bonding d. resonance e. van der Waals forces (London
dispersion forces)
172. Is used to explain the fact that the carbon-to-carbon bonds in benzene, C6H6, are identical
a. hydrogen bonding b. hybridization c. ionic bonding d. resonance e. van der Waals forces (London
dispersion forces)
173. The weight of H2SO4 (molecular weight 98.1) in 50.0 milliliters of a 6.00-molar solution is
a. 3.10 grams b. 12.0 grams c. 29.4 grams d. 294 grams e. 300. grams