Name:
Chemistry 1A – Final Review – Multiple Choice
1. 1.00 cm is equal to how many meters?
0.0100
2. 1.00 cm is equal to how many inches? (1 inch = 2.54 cm)
0.394
3. 4.50 ft is how many centimeters?
137
4. The number 0.0048 contains how many significant digits?
2
5. Express 0.00382 in scientific notation.
3.82  10-3
6. 42C is equivalent to (F = 1.8  C + 32)
108F
7. 267F is equivalent to
404 K
8. An object has a mass of 62 g and volume of 4.6 mL. Its density is
13 g/mL
9. The mass of a block is 9.43 g and its density is 2.35 g/mL. The block’s volume is:
4.01 mL
10. The mass of a piece of copper that has a volume of 9.5 mL is (dcopper = 8.92 g/mL)
85 g
11. An empty graduate cylinder has a mass of 54.772 g. When filled with 50.0 mL of an
unknown liquid, it has a mass of 101.074 g. The density of the liquid is:
0.926 g/mL
12. The conversion factor to change grams to milligrams is
0.001 g
13. If an object has a density greater than 1.00 g/mL it will sink in water.
14. A gold alloy has a density of 12.42 g/mL and contains 75.0% gold by mass. The volume
of the alloy that can be made from 255 g of pure gold is
27.4 mL
15. A lead cylinder (V = r2h) has a radius 12.0 cm and length 44.0 cm and a density of 11.4
g/mL. The mass of the cylinder is
2.27  105 g
16. The following units can be used for density except
a. g/cm3
b. kg/m3
c. g/L
d. kg/m2
17. 37.4 cm  2.2 cm equals
82 cm2
18. The following elements are among the five most abundant by mass in the Earth’s crust,
seawater, and atmosphere except
a. oxygen
b. hydrogen
c. silicon
d. aluminum
19. Which of the following is a compound?
a. lead
b. wood
c. potassium
d. water
20. Which of the following is a mixture?
a. water
b. chromium
c. wood
d. sulfur
21. How many atoms are represented in the formula Na2CrO4?
a. 3
b. 5
c. 7
d. 8
22. Which of the following is a characteristic of metals?
a. ductile
b. easily shattered c. extremely strong d. dull
23. Which of the following is a characteristic of nonmetals?
a. always a gas
c. shiny
b. poor conductor of electricity
d. combines only with metals
24. When a pure substance was analyzed, it was found to contain carbon and chlorine. The
substance must be classified as
a. an element
b. a mixture
c. a compound
d. both mixture
and compound
25. Chromium, fluorine, and magnesium have the symbols
Cr, F, Mg
26. Sodium, carbon, and sulfur have the symbols
Na, C, S
27. Coffee is an example of
homogeneous mixture
28. The number of oxygen atoms in Al(CH3COO)3 is
6
29. Which of the following is a mixture?
a. water
b. iron (II) oxide
c. sugar solution
d. iodine
30. Which is the most compact state of matter?
solid
31. Which is not a characteristic of a solution?
a. a homogeneous mixture
c. contains two or more substances
b. a heterogeneous mixture
d. has a variable composition
32. The name for the state change from solid to gas is:
sublimation
33. The number of nonmetal atoms in Al2(SO3)3 is
a. 5
b. 7
c. 12
d. 14
34. Which of the following is not a physical property?
a. boiling point
b. physical state
c. bleaching action d. color
35. Which of the following is a physical change?
a. a piece of sulfur is burned
c. a rubber band is stretched
b. a firecracker explodes
d. a nail rusts
36. Which of the following is a chemical change?
a. water evaporates
c. rocks are ground into sand
b. ice melts
d. a penny tarnishes
37. When 9.44 g of calcium are heated in air, 13.22 g of calcium oxide are formed. The
percent by mass of oxygen in the compound is
28.6%
38. Barium iodide, BaI2, contains 35.1% barium by mass. An 8.50 g sample of barium iodide
contains what mass of iodine?
5.52 g
39. Mercury (II) sulfide, HgS, contains 86.2% mercury by mass. The mass of HgS that can
be made from 30.0 g of mercury is
34.8 g
40. The changing of liquid water to ice is known as a
a. chemical change
c. homogeneous change
b. heterogeneous change
d. physical change
41. Which of the following does not represent a chemical change?
a. heating of copper in air
c. cooling of red-hot iron
b. combustion of gasoline
d. digestion of food
42. Heating 30. g of water from 20.C to 50.C requires
3.8  103 or 3800 J
43. The specific heat of aluminum is 0.900 J/gC. How many joules of energy are required
to raise the temperature of 20.0 g of Al from 10.0C to 15.0C?
90. J
44. A 100. g iron ball (specific heat = 0.473 J/gC) is heated to 125C and is placed in a
calorimeter holding 200. g of water at 25.0C. What will be the highest temperature
reached by the water?
a. 43.7C
b. 30.4C
c. 65.3C
d. 35.4C
45. Which has the highest specific heat?
a. ice
b. lead
c. water
d. aluminum
46. When 20.0 g of mercury is heated from 10.0C to 20.0C, 27.6 J of energy are absorbed.
What is the specific heat of mercury?
0.138 J/gC
47. Changing hydrogen and oxygen into water is a
a. physical change
c. conservation reaction
b. chemical change
d. nuclear reaction
48. The concept of a positive charge and a small, “heavy” nucleus surrounded by electrons
was the contribution of
Ernest Rutherford
49. The modern atomic theory was first proposed by
John Dalton
50. Whose cathode ray tube experiments were used to discover the electron?
J.J. Thomson
51. How many electrons are in an atom of
18
40
18
52. The number of neutrons in an atom of
83
139
56
Ar ?
Ba is
53. An atom of atomic number 53 and mass number 127 contains how many neutrons?
74
54. Each atom of a specific element has the same number of protons.
55. Which pair of symbols represents isotopes?
23
Na and 23
a. 10
b. 73 Li and 83 Li
c.
11 Na
63
29
Cu and
29
63
Cu
d.
12
23
Mg and
12
24
Mg
56. Two naturally occurring isotopes of an element have names and abundances as follows:
54.00 amu (20.00%) and 56.00 amu (80.00%). What is the relative atomic mass of the
element?
55.60
57. Substance X has 13 protons, 14 neutrons, and 10 electrons. Determine its identity.
3+
c. 27
13 Al
58. The mass of a chlorine atom is 5.90  10-23 g. How many atoms are in a 42.0 g sample of
chlorine?
7.12  1023
59. The number of neutrons in an atom of
61
108
47
Ag is
60. The number of electrons in an atom of
13
27
13
Al is
61. The number of protons in an atom of
30
65
30
Zn is
62. The number of protons in the nucleus of an atom of
12
24
12
Mg is
63. 4.0 g of oxygen contains
a. 1.5  1023 atoms of oxygen
b. 4.0 molar masses of oxygen
c. 0.50 mol of oxygen
d. 6.022  1023 atoms of oxygen
64. One mole of hydrogen atoms contains
a. 2.0 g of hydrogen
b. 6.022  1023 atoms of hydrogen
c. 1 atom of hydrogen
d. 12 g of carbron-12
65. The mass of one atom of magnesium is
4.037  10-23 g
66. Avogadro’s number of magnesium atoms
a. has a mass of 1.0 g
b. has a mass of 12.0 g
c. has the same mass as Avogadro’s number of sulfur atoms
d. is 1 mol of magnesium atoms
67. Which of the following contains the largest number of moles?
a. 1.0 g Li
b. 1.0 g Na
c. 1.0 g Al
d. 1.0 g Ag
68. How many moles of aluminum hydroxide are in one antacid tablet containing 400. mg of
Al(OH)3?
5.13  10-3
69. How many grams of Au2S can be obtained from 1.17 mol of Au?
249 g
70. The molar mass of Ba(NO3)2 is
261.3 g/mol
71. A 16 g sample of O2
a. is 1 mol of O2
b. contains 6.022  1023 molecules of O2
c. is 0.50 molecule of O2
d. is 0.50 molar mass of O2
72. What is the percent composition for a compound formed from 8.15 g of zinc and 2.00 g
of oxygen?
80.3% Zn, 19.7% O
73. Which of these compounds contains the largest percentage of oxygen?
a. SO2
b. SO3
c. N2O3
d. N2O5
74. 2.00 mol of CO2
a. has a mass of 56.0 g
b. contains 1.20  1024 molecules
c. has a mass of 44.0 g
d. contains 6.00 molar masses of CO2
75. In Ag2CO3, the percent by mass of
a. C is 43.5%
b. Ag is 64.2%
c. O is 17.4%
d. O is 21.9%
76. The empirical formula of the compound containing 31.0% Ti and 69.0% Cl is
TiCl3
77. A compound contains 54.3% C, 5.6% H, and 40.1% Cl. The empirical formula is
C4H5Cl
78. A compound (molar mass = 60.0g) contains 40.0% C, 6.7% H, and 53.3% O. The
molecular formula is
C2H4O2
79. How many chlorine atoms are in 4.0 mol of PCl3?
7.2  1024
80. What is the mass of 4.53 mol of Na2SO4?
644 g
81. The percent composition of Mg3N2 is
72.2% Mg, 27.8% N
82. How many grams of oxygen are contained in 0.500 mol of Na2SO4?
32.0 g
83. The empirical formula of a compound is CH. If the mass of the compound is 78.11 g,
then the molecular formula is
C6H6
84. The reaction BaCl2 + (NH4)2SO4  BaSO4 + 2 NH4Cl is an example of
double displacement
85. Balance:
4 Al + 3 O2  2 Al2O3
86. Which equation is incorrectly balanced?
a. 2 KNO3  2 KNO2 + O2
b. H2O2  H2O + O2
c. 2 Na2O2 + 2 H2O  4 NaOH + O2
d. 2 H2O  2 H2 + O2
87. The reaction 2 Al + 3 Br2  2 AlBr3 is an example of synthesis.
88. Balance:
2 PbO2  2 PbO + O2
89. Balance:
Cr2S3 + 6 HCl  2 CrCl3 + 3 H2S
90. Balance:
2 F2 + 2 H2O  4 HF + O2
91. Complete and balance:
Complete and balance:
2 NH4OH + H2SO4  (NH4)2SO4 + 2 H2O
5 H2 + V2O5  2 V + 5 H2O
92. Balance:
2 Al(OH)3 + 3 H2SO4  Al2(SO4)3 + 6 H2O
93. Balance:
2 H3PO4 + 3 Ca(OH)2  6 H2O + Ca3(PO4)2
94. Complete and balance:
Fe2(SO4)3 + 3 Ba(OH)2  3 BaSO4 + 2 Fe(OH)3
95. For the reaction 2 H2 + O2  2 H2O + 572.4 kJ which of the following is not true?
a. The reaction is exothermic.
b. 572.4 kJ of heat are liberated for each mole of water formed.
c. 2 mol of hydrogen react with 1 mol of oxygen.
d. 572.4 kJ of heat are liberated for each 2 mol of hydrogen reacted.
96. How many moles is 20.0 g of Na2CO3?
0.189 mol
97. What is the mass of 0.30 mol of BaSO4?
70. g
98. How many molecules are in 5.8 g of acetone, C3H6O?
6.0  1022
Problems 99 – 105 refer to the reaction 2 C2H4 + 6 O2  4 CO2 + 4 H2O
99. If 6.0 mol of CO2 are produced, how many moles of O2 were reacted?
9.0 mol
100.
How many moles of O2 are required for the complete reaction of 45 g of C2H4?
4.8 mol
101.
If 18.0 g of CO2 are produced, how many grams of H2O are produced?
7.37 g
102.
How many moles of CO2 can be produced by the reaction of 5.0 mol of C2H4 and
12.0 mol of O2?
8.0 mol
103.
How many moles of CO2 can be produced by the reaction of 0.480 mol of C2H4 and
1.08 mol of O2?
0.720 mol
104.
How many grams of CO2 can be produced from 2.0 g of C2H4 and 5.0 g of O2?
4.6 g
105.
If 14.0 g of C2H4 is reacted and the actual yield of H2O is 7.84 g, the percent yield in
the reaction is
43.6%
Problem 107 – 109 refer to the equation H3PO4 + MgCO3  Mg3(PO4)2 + CO2 + H2O
106.
Balance the equation. 2 H3PO4 + 3 MgCO3  Mg3(PO4)2 + 3 CO2 + 3 H2O
107.
If 20.0 g of carbon dioxide is produced, the number of moles of magnesium carbonate
used is
0.454 mol
108.
If 50.0 g of magnesium carbonate reacts completely with H3PO4, the number of
grams of carbon dioxide produced is
26.1 g
109.
When 10.0 g of MgCl2 and 10.0 g of Na2CO3 are reacted in MgCl2 + Na2CO3 
MgCO3 + 2 NaCl the limiting reactant is
Na2CO3
110.
When 50.0 g of copper is reacted with silver nitrate solution in Cu + 2 AgNO3 
Cu(NO3)2 + 2 Ag 148 g of silver is obtained. What is the percent yield of silver
obtained?
87.1%
111.
The concept of electrons existing in specific orbits around the nucleus was the
contribution of
Bohr
112.
The correct electron configuration for a fluorine atom is 1s22s22p5
113.
The correct electron configuration for 48Cd is 1s22s22p63s23p64s23d104p65s24d10
114.
The correct electron configuration of 23V is [Ar]4s23d3
115.
The number of orbitals in a d sublevel is 5
117.
The number of electrons in the third principal energy level in an atom having the
electron configuration 1s22s22p63s23p2
4
118.
The total number of orbitals that contain at least one electron in an atom having the
configuration 1s22s22p63s23p2 is 8
119.
Which of these elements has two s and six p electrons in its outer energy level?
Ar
120.
Which element is not a noble gas?
Ra
121.
Which element has the largest number of unpaired electrons?
N
122.
How many unpaired electrons are in the electron configuration [Ar]4s13d5?
6
123.
Groups 13-18 form the area on the periodic table where the electron sublevels being
filled are s and p sublevels
124.
Which of the following is an incorrect formula?
a. NaCl
b. K2O
c. AlO
125.
The lanthanide and actinide series of elements are
d. BaO
a. all artificially made
b. transition elements
c. filling in d level electrons
d. filling in f level electrons
126.
The element having the structure 1s22s22p63s23p2 is in Group 14
127.
In Group 15, the element having the smallest atomic radius is N
128.
In Group 14, the most metallic element is Sn
129.
Which group in the periodic table contains the least reactive elements?
Noble Gases
130.
Which group on the periodic table contains the Alkali Metals? 1
131.
An atom of fluorine is smaller than an atom of oxygen. One possible explanation is
that, compared to oxygen, fluorine has
a. a larger mass number
c. a greater nuclear charge
b. a smaller atomic number
d. more unpaired electrons
132.
If the size of the fluorine atom is compared to the size of the fluoride ion,
a. they would both be the same size
c. the ion is larger than the atom
b. the atom is larger than the ion
d. the sizes depend on the reaction
133.
Sodium is a very active metal because
a. it has a low ionization energy
b. it has only one outermost electron
134.
c. it has a relatively small atomic mass
d. all of the above
Which of the following formulas is incorrect?
a. Na+
b. Sc. Al3+
d. F-
135.
Which of the following molecules does not have a polar covalent bond?
a. CH4
b. H2O
c. CH3OH
d. Cl2
136.
Which of the following molecules has a dipole?
a. HBr
b. CH4
c. H2
d. CO2
Which of the following has bonding that is ionic?
a. H2
b. MgF2
c. H2O
d. CH4
137.
138.
Which of the following is a correct Lewis structure?
b. CCl4
139.
Which of the following is an incorrect Lewis structure?
a. NH2
140.
When a magnesium atom participates in a chemical reaction, it is most likely to
lose 2 electrons
141.
If X represents an element of Group 13, what is the general formula for its oxide?
X2O3
142.
Which of the following has the same electron configuration as an argon atom?
a. Ca2+
b. Cl0
c. Na+
d. K0
143.
As the difference in electronegativity between two elements decreases, the tendency
for the elements to form a covalent bond increases
144.
The number of electrons in a triple bond is 6
145.
The number of unbonded pairs of electrons in H2O is 2
146.
Which of the following does not have a noble gas configuration?
a. Na
b. Sc3+
c. Ar
d. O2-
147.
Which of these statements is not one of the principal assumptions of the kineticmolecular theory for an ideal gas?
a. All collisions of gaseous molecules are perfectly elastic.
b. A mole of any gas occupies 22.4 L at STP.
c. Gas molecules have no attraction to one another.
d. The average kinetic energy for molecules is the same for all gases at the same
temperature.
148.
Which of the following is not equal to 1.00 atm?
a. 760. cm Hg
b. 101.35 kPa
c. 760. mm Hg
d. 760. torr
149.
If the pressure on 45 mL of gas is changed from 600. torr to 800. torr, the new
volume will be 34 mL
150.
The volume of a gas is 300. mL at 740. torr and 25°C. If the pressure remains
constant and the temperature is raised to 100°C, the new volume will be 376 mL
151.
The volume of a dry gas is 4.00 L at 15.0°C and 745 torr. What volume will the gas
occupy at 40.0°C and 700. torr? 4.63 L
152.
A sample of Cl2 occupies 8.50 L at 80.0°C and 740. mm Hg. What volume will the
Cl2 occupy at STP? 6.40 L
153.
What volume will 8.00 g O2 occupy at 45.0°C and 2.00 atm?
3.26 L
154.
Measured at 65°C and 500. torr, the mass of 3.21 L of a gas is 3.5 g. The molar mass
of this gas is 46 g/mol
155.
Box A contains O2 at a pressure of 200 torr. Box B, which is identical to Box in
volume, contains twice as many molecules of CH4 as the molecules of O2 in Box A.
The temperatures of the gases are equal. The pressure in Box B is 400 torr
156.
How many liters of NO2 (at STP) can be produced from 25.0 g of Cu reacting with
concentrated nitric acid? Cu + 4 HNO3  Cu(NO3)2 + 2 H2O + 2 NO2
a. 4.41 L
b. 8.82 L
c. 17.6 L
d. 44.8 L
157.
How many liters of butane (C4H10) are required to produce 2.0 L CO2 at STP?
2 C4H10 + 13 O2  8 CO2 + 10 H2O
0.50 L
158.
What volume of CO2 (at STP) can be produced when 15.0 g C2H6 and 50.0 g O2 are
reacted? 2 C2H6 + 7 O2  4 CO2 + 6 H2O 20.0 L
159.
How many molecules are present in 0.025 mol of H2 gas? 1.5  1022
160.
Which of the following is an incorrect equation?
a. H2SO4 + 2 NaOH  Na2SO4 + 2 H2O c. 2 H2O  2 H2 + O2
b. C2H6 + O2  2 CO2 + 3 H2
d. Ca + 2 H2O  Ca(OH)2 + H2
161.
How many kilojoules are required to raise 85 g of water at 25°C to a temperature of
100°C? 27 kJ
162.
The formula for iron (II) sulfate heptahydrate is FeSO47 H2O
163.
The process by which a solid changes directly to a vapor is called sublimation
164.
The molarity of a solution containing 2.5 mol CH3COOH in 400. mL of solution is
6.3 M
165.
What volume of 0.300 M KCl will contain 15.3 g KCl?
683 mL
166.
What mass of BaCl2 will be required to prepare 200. mL of 0.150 M solution?
6.25 g
Problems 167 – 169 refer to the reaction CaCO3 + 2 HCl  CaCl2 + H2O + CO2
167.
What volume of 6.0 M HCl will be needed to react with 0.350 mol of CaCO3?
117 mL
168.
If 400. mL of 2.0 M HCl reacts with excess CaCO3, the volume of CO2 produced,
measured at STP, is 9.0 L
169.
If 5.3 g CaCl2 is produced in the reaction, what is the molarity of the HCl used if 25 mL
of it is reacted with excess CaCO3? 3.8 M
170.
How many milliliters of 6.0 M H2SO4 must you used to prepare 500. mL of 0.20 M
sulfuric acid solution?
17
171.
Which procedure is most likely to decrease the solubility of most solid in liquids?
a. stirring
c. heating the solution
b. breaking up the solid
d. increasing the pressure
172.
Which of these anions will not form a precipitate with Ag+?
a. Clb. NO3c. Br-
d. CO32-
Which of these salts are considered to be soluble in water?
a. BaSO4
b. NH4Cl
c. AgI
d. PbS
173.
174.
2 Al + 6 HCl  3 H2 + 2 AlCl3
175.
H3PO4 + 3 KOH  K3PO4 + 3 H2O
176.
6 HCl + Cr2(CO3)3  3 CO2 + 3 H2O + 2 CrCl3
177.
Which of these is not an acid?
a. H3PO4
b. H2S
c. H2SO4
d. NH3
Which of these is a nonelectrolyte?
a. CH3COOH
b. MgSO4
c. KMnO4
d. CCl4
Which of these is a weak electrolyte?
a. NaOH
b. NaCl
c. CH3COOH
d. H2SO4
178.
179.
180.
A solution has an H+ cocnentration of 3.4  10-5 M. The pH is 4.47
181.
A solution with a pH of 5.85 has an H+ concentration of 1.4  10-6 M
182.
If 16.55 mL of 0.844 M NaOH is required to titrate 10.00 mL of a hydrochloric acid
solution, the molarity of the acid solution is 1.40 M
183.
What volume of 0.462 M NaOH is required to neutralize 20.00 mL of 0.391 M HNO3?
16.9 mL
184.
25.00 mL of H2SO4 solution requires 18.92 mL of 0.1024 M NaOH for complete
neutralization. The molarity of the acid is 0.03875 M
185.
What is the pH of a 0.00015 M HCl solution?
186.
The chloride ion concentration in 300. mL of 0.10 M AlCl3 is 0.30 M
187.
The amount of BaSO4 that will precipitate when 100. mL of 0.10 M BaCl2 and 100. mL
of 0.10 M Na2SO4 are mixed is 0.010 mol
188.
The equation CH3COOH + H2O ‡ˆ ˆ †ˆ H3O+ + CH3COO- implies that
a. If you start with 1.0 mol CH3COOH, 1.0 mol H3O+ and 1.0 mol CH3COO- will be
produced.
b. An equilibrium exists between the forward reaction and the reverse reaction.
c. At equilibrium, equal molar amounts of all four substances will exist.
d. The reaction proceeds all the way to the products, then reverses, going all the way
back to the reactants.
189.
If the reaction A + B ‡ˆ ˆ †ˆ C + D is initially at equilibrium, and then more A is added,
which of the following is not true?
a. More collisions of A and B will occur; the rate of the forward reaction will thus be
increased.
b. The equilibrium will shift towards the right.
c. The moles of B will be increased.
d. The moles of D will be increased.
190.
In the equilibrium reaction N2 + 2 O2 ‡ˆ ˆ †ˆ 2 NO2 as the pressure is increased, the
amount of NO2 formed
a. increases
c. remains the same
b. decreases
d. increases and decreases irregularly
191.
If
238
92
U loses an alpha particle, the resulting nuclide is
192.
If
210
82
Pb loses a beta particle, the resulting nuclide is 210
83 Bi
193.
In the equation
194.
In the nuclear equation
195.
If you started with 40 g of the isotope
after 9.0 minutes? 5.0 g
209
83
Bi + ? 
45
21
210
84
3.82
234
90
Th
Po + 01 n , the missing particle would be 21 H 
Sc + 01 n  X + 11 H the nuclide X that is formed is
210
84
45
20
Ca
Po (t1/2 = 3.0 min), how much would be left
Consider the nuclides iodine-131, radon-222, and uranium-238.
a. How many protons and neutrons do these nuclides contain?
iodine-131: 53 p+/78 n
radon-222: 86 p+/136 n uranium-238: 92 p+/146 n
b. How many nucleons do these nuclides contain?
iodine-131: 131
radon-222: 222
uranium-238: 238
c. Write the equation for the alpha and beta decays for these nuclides.
131
53
I 
 24 a  127
51 Sb
131
53
I 
 10  131
54 Xe
222
86
Rn 
 24 a  218
84 Po
222
86
Rn 
 10  222
87 Fr
U 
 24 a  234
90Th
238
92
U 
 10  238
93 Np
238
92
d. If you have 3.00 g of these samples, how much remains after 50.0 days?
iodine-131 = 0.0398 g
radon-222 = 3.44  10-4 g
uranium-238 = 3.00 g
Briefly describe the contributions of Rutherford, Bohr, Thomson , and Schrödinger.
Thomson – discovered the electron using the cathode ray tube, proposed “plum
pudding model” – electrons were interspersed in positively charged space.
Rutherford – discovered the positively charged nucleus in the gold foil experiment;
proposed electrons orbited nucleus.
Bohr – used the idea of orbiting electrons the nucleus in energy levels to explain how
light was produced.
Schrödinger – postulated the idea of orbitals – regions of space where electrons were
most likely to be found.
Briefly summarize the concepts of the law of conservation of mass, law of definite
proportions, and law of multiple proportions.
Law of Conservation of Mass – in any non-nuclear (chemical) reaction, the mass of
reactants and products must be equivalent.
Law of Definite Proportions – all molecules of the same compound have the same
integer ratios of elements to one another.
Law of Multiple Proportions – elements can combine in different integer ratios to
form different compounds.
Construct an energy level diagram and label the following energy levels: (E1 = 0 J; E2 =
3.44  10-19 J; E3 = 5.66  10-19 J; E4 = 6.79  10-19 J; E5 = 7.16  10-19 J; E6 = 7.44  1019
J; E7 = 7.57  10-19 J;  = 7.68  10-19 J).
Complete the table below:
Transition
4th to 6th
6th to 3rd
5th to ground
state
7th to 4th
2nd to 5th
eject e- from
E2
7th to 5th
6th to 2nd
Energy (J)
Frequency
(Hz)
Wavelength
(m)
Energy (eV) E absorbed
or released?
6.50E-20
1.78E-19
9.81E+13
2.69E+14
3.06E-06
1.12E-06
0.406
1.11
absorbed
released
7.16E-19
7.80E-20
3.72E-19
1.08E+15
1.18E+14
5.61E+14
2.77E-07
2.55E-06
5.34E-07
4.47
0.487
2.32
released
released
absorbed
4.24E-19
4.10E-20
4.00E-19
6.40E+14
6.19E+13
6.04E+14
4.69E-07
4.85E-06
4.97E-07
2.65
0.256
2.50
absorbed
released
released
Give an element that belongs to each of the following categories:
a. Halogens
F,Cl, Br, I, At
b. Lanthanide Series
Ce - Lu
c. Alkali Metals
Li, Na, K, Rb, Cs, Fr
d. Transition Metals
Group 3 – 12
e. Noble Gases
He, Ne, Ar, Kr, Xe, Rn
f. Alkali Earth Metals
Be, Mg, Ca, Sr, Ra
g. Period 5
Rb – Xe
h. Chalcogens
O, S, Se, Te, Po
i. p-block elements
Group 13 – 18 not He
Give the element that fits the following criteria:
a. largest in Group 14
b. highest electron affinity in period 2
c. lowest 1st ionization energy in group 8
d. smallest in period 5
e. least reactive of the halogens
f. lowest electron affinity in the chalcogens
g. greatest 1st ionization energy in period 3
Pb
F
Hs
Xe
At
Po
Cl
Draw electron dot diagrams for the following elements:
a. iodine
b. calcium
c. phosphorous
d. xenon
e. aluminum
Draw the following molecules and label their dipoles. Determine if the molecule has a net
dipole or not.
a. CCl4
no net dipole
b. Li3N
net dipole
unshared pr
electrons
c. SiFClBr2
d. CaCl2
net dipole
no net dipole
diff elems on Si
e. K2Se
net dipole
unshared
pr e- on Se
In which of the following is the formula correct for the name given?
a. copper (II) sulfate, CuSO4
b. ammonium hydroxide, NH4OH
c. mercury (I) carbonate, HgCO3
d. phosphorous triiodide, PI3
e. calcium acetate, Ca(CH3COO)2
f. hypochlorous acid, HClO
g. dichlorine heptoxide, Cl2O7
h. magnesium iodide, MgI
i. sulfurous acid, H2SO3
j. potassium permanganate, KMnO4
k. lead (II) chromate, PbCrO4
l. iron (II) phosphate, FePO4
m. mercury (II) sulfate, HgSO4
n. dinitrogen pentoxide, N2O5
o. sodium hypochlorite, NaClO
p. sodium dichromate, Na2Cr2O7
q. cadmium cyanide, Cd(CN)2
r. bismuth (III) oxide, Bi3O2
s. carbonic acid, H2CO3
t. silver oxide, Ag2O
u. iron (III) iodide, FeI2
v. tin (II) fluoride, TiF2
w. carbon monoxide, CO
x. phosphoric acid, H3PO3
y. sodium bromate, Na2BrO3
z. hydrosulfuric acid, H2S
aa. potassium hydroxide, POH
ab. sodium carbonate, Na2CO3
ac. zinc sulfate, ZnSO4
ad. sulfur trioxide, SO3
ae. tin (IV) nitrate, Sn(NO3)4
af. iron (II) sulfate, FeSO4
ag. chloric acid, HCl
ah. aluminum sulfide, Al2S3
ai. cobalt (II) chloride, CoCl2
aj. acetic acid, CH3COOH
ak. zinc oxide, ZnO2
al. tin (IV) nitrite, Sn(NO3)4
Which of the six types of reactions are represented by the following?
C – Combustion
S – Synthesis
SD – Single Displacement
C2H6 + O2 → CO2 + H2O
C, R
P4O6 + H2O → H3PO3
S
KClO3 → KCl + KClO4
D, R
C3H7O2 → C + H2 + O2
D, R
NBr3 + NaOH → N2 + NaBr + HBrO R
H2SO4 + KOH → K2SO4 + H2O DD
C3H8 + O2 → CO + H2O
C, R
CaO + SiO2 → CaSiO3
S
Mg(NO3)2 + Ca → Ca(NO3)2 + Mg SD, R
HI + F2 → HF + I2
SD, R
H2CO3 → H2O + CO2
D
NaOH + H3PO4 → Na3PO4 + H2O DD
C6H12 + O2 → CO2 + H2O
C, R
Ca3(PO4)2 + C → Ca3P2 + CO
R
H2S + O2 → SO2 + H2O
R
N2O5 → NO2 + O2
D, R
Rb + Br2 → RbBr
D, R
S + O2 → SO2
S, R
CH4 + O2 → CO2 + H2O
C, R
AgNO3 + HCl → AgCl + HNO3 DD
HCl + K2CO3 → KCl + CO2 + H2O DD, D
MnSO4 + K → K2SO4 + Mn SD, R
D – Decomposition
R – Redox
DD – Double Displacement
Which elements are oxidized and reduced in the following reactions?
ox – Xe red – F
ox – O
red – Fe
ox – N (NH3) red – N (NO)
ox – H
red – Fe
ox – C
red – Fe
ox – Hg (HgO) red – Hg (Hg)
Predict the products for the following reactions:
H2 + I2 → HI
HClO3 + NaOH → H2O + NaClO3
MgCl2 + K2CO3 → KCl + MgCO3
HOCH2CH2CH2OH + O2

 CO2 + H2O

Mg + H2SO4 → MgSO4 + H2
Zn(NO3)2 + BaCl2 → Ba(NO3)2 + ZnCl2
low oxygen level
 CO + H2O
C4H10 + O2 

Al + F2 → AlF3
H2O → H2 + O2
Sr(NO3)2 + KOH → Sr(OH)2 + KNO3
MgCl2 + K2CO3 → MgClO3 + KCl
HBr + Na2CO3 → NaBr + CO2 + H2O
CH3CH2CH2OH + O2 
 CO2 + H2O

Li + ZnSO4 → Zn + Li2SO4
ZnSO4 + BaCl2 → BaSO4 + ZnCl2
Determine which of the following reactions work or not:
Al  ZnBr2 
 Zn  AlBr3
AlCl3  Zn(CH 3COO )2 
 NR
Ca  HNO3 
 H 2  Ca ( NO3 ) 2
CaCO3  HNO3 
 H 2O  CO2  Ca ( NO3 ) 2
Pb  CuCl2 
 PbCl2  Cu
Pb(CH 3COO ) 2  CuCl2 
 NR
Li  K 2 SO4 
 K  Li2 SO4
Ca (OH ) 2  H 2 SO4 
 HOH  CaSO4
F2  ZnBr2 
 Br2  ZnF2
SnCl2  ZnBr2 
 NR
Ba  Fe( NO3 )3 
 Fe  Ba ( NO3 ) 2
Ba( NO3 ) 2  FeSO4 
 Fe( NO3 ) 2  BaSO4
I 2  NaCl 
 NR
KI  NaCl 
 NR
Mg  K 3 PO4 
 NR
MgSO4  K 3 PO4 
 K 2 SO4  Mg3 ( PO4 )2
Use thermodynamic data to determine if the following reactions work at the given
temperature:
C3H8 (g) + 5 O2 (ℓ) → 3 CO2 (g) + 4 H2O (g) at 500°C
H = -2043.9 kJ/mol
S = 100.3 J/mol K
G = -2121.4 kJ/mol – spontaneous
2 SO2 (g) + O2 (g) → 2 SO3 (g) at -150°C

H = -197.6 kJ/mol
S = -188 J/mol K
G = -174.5 kJ/mol – spontaneous
2 Fe (s) + Al2O3 (s) → 2 Al (s) + Fe2O3 (s) at 455°C
H = 851.5 kJ/mol
S = 38.5 J/mol K
G = 823.5 kJ/mol – nonspontaneous
Interpret the meaning of positive and negative enthalpy, entropy, and free energy values.
Positive enthalpy – endothermic
Negative enthalpy – exothermic
Positive entropy – greater disorder
Negative entropy – less disorder
Positive free energy – nonspontaneous
Negative free energy – spontaneous
Answer the following stoichiometry problems:
1. a. 185.8 g
2. a. 42.37 mL
b. 890.0 L
b. 1.48 g
c. 79.87 g
Use the various concentration formulas to answer the following:
a. What will be the molarity of a solution made by dissolving 32.67 g of potassium
sulfite in 545.0 mL of water?
0.3787 M
b. A certain chemical process requires 25.00 L of a 0.766 M solution of cesium chloride
What mass of c cesium chloride is needed? 3220 g
c. What is the volume of a 0.553 M calcium chloride solution that contains 1.77g of
solute?
28.8 mL
d. Write its dissociation equation for iron (III) bromide.
FeBr3 → Fe3+ + 3 Br-
If 41.5 g of solute is dissolved to a final volume of 2.45 L, what is the molarity of the
solution? 0.0573 M
Determine the concentration of each ion in solution. Fe3+ = 0.0573 M
3 Br- = 0.172 M
e. Write its dissociation equation for sulfuric acid. H2SO4 → 2 H+ + SO42If 1.55  10-4 mol of acid are dissolved in 2.55 L of solution, what is the molarity of
the acid?
6.07 × 10-5 M
What will the pH of the solution? What is the pOH of the solution?
pOH = 10.084
pH = 3.916
Use the various gas laws to answer the following:
a. What will be the pressure (in mbar) of 23.55 g of fluorine in an 8.68 L container at
124.7C?
2360 mbar
b. Determine the volume of a hydrogen balloon that is originally at 1.51 L in a room at
23.0C and is moved outside where the temperature is 5.05C. 1.41 L
c. A glass container holds a mixture of gases A, B, and C. The total pressure of the
gases is 4.00 atm. If there is twice as much B as C, and twice as much A as B, then
what is the pressure of each gas? PA = 2.29 atm
PB = 1.14 atm PC = 0.571 atm
d. At constant pressure and temperature, 2.33 mol of xenon occupies 52.16 L. If
conditions do not change, what will be the volume of 350.0 g of xenon? 59.5 L
e. What will be the pressure of a gas that is enclosed in a 4.33 L container at 348 kPa
and is moved to a container that is 1.22 L? 1240 kPa
f. 45.0 g of a gas is held at 225 K inside a 91.8 L container. The pressure inside the
container is 0.566 atm. Could the gas be methane, CH4? Yes, 2.81 mol
g. Find the density of oxygen held at 2.33 atm and a temperature of 175 K.
5.19 g/L
Consider the reaction and the stresses applied to it:
85 kJ + A (g) + 2 B (s) ‡ˆ ˆ †ˆ C (g) + 2 D (g)
Equilibrium shift
Increase heat
right
A
↓
B
↓
C
↑
D
↑
Decrease volume
left
↑
↑
↓
↓
Addition of C
left
↑
↑
↑
↓
Removal of D
right
↓
↓
↑
↓
Decrease heat
left
↑
↑
↓
↓
Increase pressure
left
↑
↑
↓
↓