Download Practice Multiple Choice Questions for the Chemistry Final Exam

Practice Multiple Choice Questions for the Chemistry Final Exam 2012
1.
One chemical property of matter is
a) boiling point.
c) reactivity.
b) texture.
d) density.
2.
The state of matter in which a material has definite shape and definite volume is the
a) liquid state.
b) solid state.
c) gaseous state.
d) vaporous state.
3.
Under ordinary conditions of temperature and pressure, the particles in a gas are
a) closely packed.
b) very far from each other.
c) held in fixed positions.
d) able to slide past each other.
4.
A horizontal row of elements in the periodic table is called a(n)
a) group.
b) period.
c) family.
d) octet.
5.
Elements in a group in the periodic table can be expected to have similar
a) atomic masses.
b) atomic numbers.
c) numbers of neutrons.
d) properties.
6.
The symbol that represents the measured unit for volume is
a) mL.
b) mg.
c) mm.
d) cm.
7.
A measure of the quantity of matter is
a) density.
c) volume.
8.
9.
10.
b) weight.
d) mass.
To determine density, the quantities that must be measured are
a) mass and weight.
b) volume and weight.
c) volume and concentration
d) volume and mass.
The density of aluminum is 2.70 g/m3. The volume of a solid piece of aluminum is 1.50 cm3.
Find its mass.
a) 1.50 g
b) 1.80 g
c) 2.70 g
d) 4.05 g
1.06 L of water is equivalent to
a) 0.00106 mL.
c) 106 mL.
11. Convert -25'C to the kelvin scale.
a) -323. K
c) 248. K
b) 10.6 mL.
d) 1060 mL.
b) -248. K
d) 323. K
12. In oxides of nitrogen, such as N2O, NO, NO2, and N2O3, atoms combine in small whole-number
ratios. This evidence supports the law of
a) conservation of mass.
b) multiple proportion.
c) definite composition.
d) mass action.
1
13.
Who was the schoolmaster who studied chemistry and proposed an atomic theory?
a) John Dalton
b) Jons Berzehus
c) Robert Brown
d) Dmitri Mendeleev
14.
According to Dalton’s atomic theory, atoms
a) of different elements
b) can be divided into
combine in simple
protons, neutrons, and
whole-number ratios to form
electrons.
compounds.
c) of all elements are identical d) can be destroyed in
in size and mass.
chemical reactions.
15.
The discovery of the electron resulted from experiments using
a) gold foil.
b) cathode rays.
c) neutrons.
d) alpha particles.
16.
Who discovered the nucleus by bombarding gold foil with positively charged particles and
noting that some particles were widely deflected?
a) Rutherford
b) Dalton
c) Chadwick
d) Bohr
17. Rutherford fired positively charged particles at metal foil and concluded that most of the mass
of an atom was
a) in the electrons.
b) concentrated in the nucleus.
c) evenly spread throughout d) in rings around the atom.
the atom
18. A nuclear particle that has about the same mass as a proton, but with no electrical charge, is
called a(n)
a) nuclide.
b) neutron.
c) electron.
d) isotope.
19. An atom is electrically neutral because
a) neutrons balance the
b) nuclear forces stabilize the
protons and electrons.
charges.
c) the numbers of protons and d) the numbers of protons and
electrons are equal
neutrons are equal.
20. Atoms of the same element that have different masses are called
a) moles.
b) isotopes.
c) nuclides.
d) neutrons.
21. All atoms of the same element have the same
a) atomic mass.
b) number of neutrons.
c) mass number.
d) atomic number.
22. An aluminum isotope consists of 13 protons, 13 electrons, and 14 neutrons. Its mass
number is
a) 13.
b) 14.
c) 27.
d) 40.
2
23. Carbon-14 (atomic number 6), the radioactive nuclide used in dating fossils, has
a) 6 neutrons.
b) 8 neutrons.
c) 10 neutrons.
d) 14 neutrons.
24.The number of atoms in 1 mol of carbon is
a) 6.02 x 1022.
b) 6.02 x 1023.
22
c) 5.02 x 10 .
d) 5.02 x 1023.
25. Molar mass
a) is the mass in gram of one b) is numerically equal to the
mole of a substance.
average atomic mass of the
element.
c) both a and b
d) neither a nor b
26. A sample of tin (atomic mass 118.69 amu) contains 3.01 x 1023 atoms. The mass of the sample is
a) 3.01
b) 59.3 g.
c) 72.6 g.
d) 11 g.
27. A bright line spectrum of an atom is caused by the energy released when electrons
a) jump to a higher energy
b) fall to a lower energy level.
level.
c) absorb energy and jump to
d) absorb energy and fall to a
a higher energy level.
lower energy level.
28. For an electron in an atom to change from the ground state to an excited state,
a) energy must be released.
b) energy must be absorbed.
c) radiation must be emitted
d) the electron must make a transition from a higher
to a lower energy level.
29. The set of orbitals that are dumbbell-shaped and directed along the x, y, and z axes are
called
a) d orbitals.
b) p orbitals.
c) f orbitals.
d) s orbitals.
30. The letter designations for the first four sublevels with the number of electrons that can be
accommodated in each sublevel are
a) s:1, p:3, d:10, and f:14.
b) s:1, p:3, d.5, and f 7.
c) s:2, p:6, d.10, and f 14.
d) s:1, p:2, d: 3, and f 4.
31. Which of the following rules requires that each of the p orbitals at a particular energy level
receive one electron before any of them can have two electrons?
a) Hund’s rule
b) the Pauli exclusion principle
c) the Aufbau principle
d) the quantum rule
32. What is the electron configuration for nitrogen, atomic number 7?
a) 1s2 2s2 2p3
b) 1s2 2s3 2p2
2
3
l
c) 1s 2s 2p
d) 1s2 2s2 2p2 3s1
33. Mendeleev predicted that the spaces in his periodic table represented
a) isotopes.
b) radioactive elements.
c) permanent gaps.
d) undiscovered elements.
3
34. The discovery of the noble gases changed Mendeleev's periodic table by adding a new
a) period.
b) series.
c) group .
d) sublevel block.
35. In the modern periodic table, elements are ordered according to
a) decreasing atomic mass.
b) Mendeleev's original design.
c) increasing atomic number. d) the date of their discovery.
36. Krypton, atomic number 36, is the fourth element in Group 18. What is the atomic number of
xenon, the fifth element in Group 18?
a) 54
b) 68
c) 72
d) 90
37. The electron configuration of aluminum, atomic number 13, is [Ne] 3s2 3pl. Aluminum is in
period
a) 2.
b) 3.
c) 6.
d) 13.
38. Calcium, atomic number 20, has the electron configuration [Ar] 4s2. In what period is calcium?
a) Period 2
b) Period 4
c) Period 8
d) Period 20
39. The energy required to remove an electron from an atom is the atom’s
a) electron affinity.
b) electron energy.
c) electronegativity.
d) ionization energy.
40. A measure of the ability of an atom in a chemical compound to attract electrons is called
a) electron affinity.
b) electron configuration
c) electronegativity.
d) ionization potential.
41. Across a period in the periodic table, atomic radii
a) gradually decrease.
b) gradually decrease, then
sharply increase.
c) gradually increase.
d) gradually increase, then
sharply decrease.
42. The number of valence electrons in Group 17 elements is
a) 7.
b) 8.
c) 17.
d) equal to the period number.
43. The electrons involved in the formation of a chemical bond are called
a) dipoles.
b) s electrons.
c) Lewis electrons.
d) valence electrons.
44. The chemical bond formed when two atoms share electrons is called a(n)
a) ionic bond.
b) orbital bond.
c) Lewis structure.
d) covalent bond.
45. If the atoms that share electrons have an unequa1 attraction for the electrons, the bond is called
a) nonpolar.
b) polar.
c) ionic.
d) dipolar.
4
46. The electron configuration of nitrogen is 1s2 2s2 2p3. How many more electrons does nitrogen
need to satisfy the octet rule?
a) 1
b) 3
c) 5
d) 8
47. A formula that shows the types and numbers of atoms combined in a single molecule is a(n)
a) molecular formula.
b) ionic formula.
c) Lewis structure.
d) covalent formula.
48. VSEPR theory is a model for predicting
a) the strength of metallic
b) the shape of molecules.
bonds.
c) lattice energy values.
d) ionization energy.
49. The following molecules contain polar bonds. The only polar molecule is
a) CCl4
b) CO2
c) NH3
d) CH4
50. How many atoms of fluorine are present in a molecule of carbon tetrafluoride, CF4?
a) 1
b) 2
c) 4
d) 5
51. The formula for carbon dioxide, CO2, can represent
a) one molecule of carbon
b) 1 mol of carbon dioxide
dioxide.
molecules.
c) one molar mass of carbon dioxide
d) all of the above.
52. What is the formula for aluminum sulfate?
a) AlSO4
b) Al2SO4
c) Al2(SO4)3
d) Al(SO4)3
53. Name the compound K2SO4.
a) potassium sulfate
c) potassium sulfide
b) potassium sulfur tetroxide
d) dipotassium sulfate
54. What is the metallic ion in copper(II) chloride?
a) Co2+
b) O22+
c) Cu
d) Cl55. Name the compound N2O4.
a) sodium tetroxide
c) nitrous oxide
b) dinitrogen tetroxide
d) binitrogen oxide
56. What is the formula for sulfur dichloride?
a) NaCl2
b) SCl2
c) S2Cl
d) S2Cl2
57. The molar mass of MgI2 is
a) the sum of the masses of
1 mol of Mg and 2 mol of I.
c) the sum of the masses of
2 mol of Mg and 2 mol of I.
b) the sum of the masses of
1 mol of Mg and I mol of I.
d) impossible to calculate.
5
58. The molar mass of NO2 is 46.01 g/mol. How many moles of NO2 are present in 114.95 g?
a) 0.4003 mol
b) 1.000 mol
c) 2.498 mol
d) 114.95 mol
59. What is the percentage composition of CF4?
a) 20% C, 80% F
b) 13.6% C, 86.4% F
c) 16.8% C, 83.2% F
d) 81% C, 19% F
60. What is the empirical formula for a compound that is 43.6% phosphorus and 56.4% oxygen?
a) P3O7
b) PO3
c) P2O3
d) P2O5
61. The molecular formula for vitamin C is C6H8O6. What is the empirical formula?
a) CHO
b) CH2O
c) C3H4O3
d) C2H4O2
62. A compound’s empirical formula is HO. If the formula mass is 34 amu, what is the molecular
formula?
a) H2O
b) H2O2
c) HO3
d) H2O3
63. In writing an equation that produces hydrogen gas, the correct representation of hydrogen gas is
a) H.
b) 2H.
c) H2.
d) OH.
64. Which equation is NOT balanced?
a) 2H2 + O2  2H2O
c) 2Mg + O2  2MgO
b) Al + 3HCl  AlCl3 + H2
d) CaCl2 + 2NaBr  CaBr2 + 2NaC
65. The reaction Mg(s) + 2HCI(aq)  H2(g) + MgCl2(aq) is a
a) composition reaction
b) decomposition reaction.
c) single-replacement reaction. d) double-replacement reaction.
66. The reaction Pb(NO3)2(aq) + 2KI(aq)  PbI2(S) + 2KNO3(aq) is a
a) double-replacement reaction. b) synthesis reaction.
c) decomposition reaction.
d) combustion reaction.
67. What is the balanced equation when aluminum reacts with copper(II) sulfate?
a) Al + Cu2S A12S + Cu
b) 2Al + 3CuSO4  Al2(SO4)3 + 3Cu
c) Al + CuSO4  AlSO4 + Cu
d) 2Al + Cu2SO4 -4 A12SO4 + 2Cu
68. In the reaction N2 + 3H22NH3, what is the mole ratio of nitrogen to ammonia?
a) 1:1
b) 1:2
c) 1:3
d) 2:3
69. In the equation 2KClO3  2KCI + 3O2, how many moles of oxygen are produced when 3.0 mol
of KClO3 decompose completely?
a) 1.0 mol
b) 2.5 mol
c) 3.0 mol
d) 4.5 mol
6
70. For the reaction 2HNO3 + Mg(OH)2  Mg(NO3)2 + 2H2O, how many grams of magnesium
nitrate are produced from 8.00 mol of nitric acid, HNO3?
a) 148 g
b) 445 g
c) 593 g
d) 818 g
71. For the reaction 3Fe + 4H2O  Fe3O4 + 4H2, how many moles of iron oxide are produced from
500 g of iron?
a) 1 mol
b) 3 mol
c) 9 mol
d) 12 mol
72. For the reaction SO3 + H2O  H2SO4, calculate the percent yield if 500. g of sulfur trioxide
react with excess water to produce 575 g of sulfuric acid.
a) 82.7%
b) 88.3%
c) 91.2%
d) 93.9%
73. An ideal gas is an imaginary gas
a) not made of particles.
b) that conforms to all of the
assumptions of the kinetic theory.
c) whose particles have zero d) made of motionless
mass.
particles.
74. Which is an example of gas diffusion?
a) inflating a flat tire
c) a cylinder of oxygen stored
b) the perfume spreading through a room
d) gas escaping from a hole in a balloon
75. According to the kinetic-molecular theory, how does a gas expand?
a) Its particles become larger. b) Collisions between particles become elastic.
c) Its temperature rises.
d) Its particles move apart in straight lines
76. Which is an example of effusion?
a) air slowly escaping from a
pinhole in a tire
c) helium dispersing into a
room after a balloon pops
b) the aroma of a cooling pie
spreading across a room
d) oxygen and gasoline fumes
mixing in an automobile
carburetor
77. What does the constant bombardment of gas molecules against the inside walls of a container
produce?
a) temperature
b) density
c) pressure
d) diffusion
78. Convert the pressure 0.75 atm to mm Hg.
a) 101.325 mm Hg
b) 430 mm Hg
c) 570 mm Hg
d) 760 mm Hg
79. Standard temperature is exactly
a) 100 C.
c) 0 C.
b) 273 C.
d) 0 K.
80. Standard pressure is exactly
a) 1 atm
c) 101.325 atm
b) 760 atm.
d) 101 atm.
7
81. Pressure and volume changes at a constant temperature can be calculated using
a) Boyle's law.
b) Charles's law.
c) Kelvin’s law.
d) Dalton's law.
83. The volume of a gas is 5.0 L when the temperature is 5.0 C. If the temperature is increased to
10.0 C without changing the pressure, what is the new volume?
a) 2.5 L
b) 4.8 L
c) 5.1 L
d) 10.0 L
84. A 150.0 L sample of gas is collected at 1.20 atm and 25.0 C. What volume does the gas have at
1.50 atm and 20.0 C?
a) 94 L
b) 118 L
c) 143 L
d) 183 L
85. In the equation H2(g) + Cl2(g)  2HCI(g), one volume of hydrogen yields how many volumes
of hydrogen chloride?
a) 1
b) 2
c) 3
d) 4
86. At constant temperature and pressure, gas volume is directly proportional to the
a) molar mass of the gas.
b) number of moles of gas.
c) density of the gas at STP. d) rate of diffusion.
87. According to Avogadro's law, 1 L of H2(g) and 1 L of O2(g) at the same temperature and
pressure
a) have the same mass.
b) have unequal volumes.
c) contain 1 mol of gas each. d) contain equal numbers of molecules.
88. The standard molar volume of a gas at STP is
a) 22.4 L.
b) g/22.4 L.
c) g-mol wt/22.4 L.
d) 1 L.
89. A 1.00 L sample of a gas has a mass of 1.25 g at STP. What is the mass of 1 mol of this gas?
a) a little less then 1.0 g
b) 1.25 g
c) 22.4 g
d) 28.0 g
90. According to Graham’s law, two gases at the same temperature and pressure will have different
rates of diffusion because they have different
a) volumes.
b) molar masses.
c) kinetic energies.
d) condensation points.
91. Why are water molecules polar?
a) They contain two kinds of b) The electrons in the covalent
atoms.
bonds spend more time closer
to the oxygen nucleus.
c) The hydrogen bonds are
d) They have covalent bonds.
weak.
92. What is the freezing point of water at standard pressure?
a) -10 C
b) 0 C
c) 4'C
d) 32 C
8
93. What is the boiling point of water at standard pressure?
a) 100 C
b) 112 C
c) 212 C
d) 200 C
94. Which of the following is a pure substance?
a) water
b) milk
c) soil
d) concrete
95. Sugar in water is an example of which solute-solvent combination?
a) gas-liquid
b) liquid-liquid
c) solid-liquid
d) liquid-solid
96. Solutions that conduct electricity are called
a) ions
b) super solutions
c) electrolytes
d) nonelectrolytes
97. If the amount of solute present in a solution at a given temperature is less than the maximum
amount that can dissolve at that temperature, the solution is said to be
a) saturated.
b) unsaturated.
c) supersaturated.
d) concentrated.
98. Which of the following is likely to produce crystals if disturbed?
a) an unsaturated solution
b) a supersaturated solution
c) a saturated solution
d) a concentrated solution
99. What is the molarity of a solution that contains 0.202 mol KC1 in 7.98 L solution?
a) 0.0132 M
b) 0.0253 M
c) 0.459 M
d) 1.363 M
100. How many moles of HCI are present in 0.70 L of a 0.33 M HCI solution?
a) 0.23 mol
b) 0.28 mol
c) 0.38 mol
d) 0.47 mol
101. An NaOH solution contains 1.90 mol of NaOH, and its concentration is 0.555 M. What is its
volume?
a) 0.623 L
b) 0.911 L
c) 1.05 L
d) 3.42 L
102. How many milliliters water are needed to make a 0.171 M solution that contains 1.00 g of
NaCl?
a) 100 mL
b) 1000 mL
c) 171 mL
d) 17.1 mL
103. How many moles of ions are produced by the dissociation of 1 mol of MgCl2?
a) 0 mol
b) 1 mol
c) 2 mol
d) 3 mol
104. Colligative properties depend on
a) the crystal size of the solute
c) the physical properties of
the solute particles.
b) the concentration of solute
d) the boiling point and
freezing point of the solution.
9
105. Compared with a 0.01 m sugar solution, a 0.01 m KCl solution has
a) the same freezing-point
b) about twice the
depression.
freezing-point depression.
c) the same freezing-point
d) about six times the
elevation.
freezing-point elevation.
106. Electrolytes affect colligative properties differently than do nonelectrolytes because
electrolytes
a) are volatile.
b) have lower boiling points.
c) produce fewer moles of
d) produce more moles of
solute particles per mole of solute particles per mole of
solute.
solute.
107. Acids taste
a) sweet.
c) bitter.
b) sour.
d) salty.
108. Acids react with
a) bases to produce salts and
water.
c) water to produce bases and
salts.
b) salts to produce bases and
water.
d) neither bases, salts, nor
water.
109. Bases taste
a) soapy.
c) sweet.
b) sour.
d) bitter.
110. A binary acid contains
a) two hydrogen atoms.
c) hydrogen and two other
elements.
b) hydrogen and one other
element.
d) hydrogen and three other
elements.
111. According to the traditional Arrhenius definition, an acid contains
a) hydrogen and does not
b) hydrogen and ionizes to
ionize.
form hydrogen ions.
c) oxygen and ionizes to form d) oxygen and ionizes to form
hydroxide ions.
oxygen ions.
112. A substance that ionizes nearly completely in aqueous solutions and produces H3O+ is a
a) weak base.
b) strong base.
c) weak acid.
d) strong acid.
113. A Bronsted-Lowry acid is
a) an electron-pair acceptor.
c) a proton acceptor.
b) an electron-pair donor.
d) a proton donor.
10
114. What is neutralization?
a) an acid-base reaction that
does not include dissocation
c) a reaction of oxygen
and hydrogen ions to
form water molecules
115. Pure water contains
a) water molecules only.
c) hydroxide ions only.
b) a reaction of hydronium ions and
hydroxide ions to form a salt
d) a reaction of an acid and a base
to form water molecules and a salt
b) hydronium ions only.
d) water molecules, hydronium
ions, and hydroxide ions.
116. What is the concentration of H3O+ in pure water?
a) 10-7 M
b) 0.7 M
c) 55.4 M
d) 107 M
117. Which expression represents the pH of a solution?
a) log[H3O+1]
b) -log[H3O+1]
c) log[OH-]
d) -log[OH-]
118. If [H3O+1] of a solution is less than [OH-1], the solution
a) is always acidic.
b) is always basic.
c) is always neutral.
d) might be acidic, basic, or neutral.
119. What is the pH of a neutral solution at 25 C?
a) 0
b) 1
c) 7
d) 14
120. The pH scale in general use ranges from
a) 0 to 1.
b) - 1 to 1.
c) 0 to 7.
d) 0 to 14.
121. The pH of an acidic solution is
a) less than 0.
c) greater than 7.
b) less than 7.
d) greater than 14.
122. The pH of a basic solution is
a) less than 0.
c) greater than 7.
b) less than 7.
d) greater than 14.
123. If [H3O+] = 1.7 x 10-3 M, what is the pH of the solution?
a) 1.81
b) 2.13
c) 2.42
d) 2.77
124. What is the pH of a 0.027 M KOH solution?
a) 6.47
b) 12.43
c) 12.92
d) 14.11
125. What is the hydronium ion concentration of a solution whose pH is 4.12?
a) 4.4 x 10-8 M
b) 5.1 X 10-6 M
c) 6.4 x 10-5 M
d) 7.6 x 10-5 M
11
126. Dyes with pH-sensitive colors are used as
a) primary standards.
b) indicators.
c) titrants.
d) None of the above
127.In an acid-base titration, equivalent quantities of hydronium ions and hydroxide ions are present
a) at the beginning point.
b) at the midpoint.
c) at the endpoint.
d) throughout the titration.
128. What is the molarity of an HCL solution if 125 mL is neutralized in a titration by 76.0 mL of
1.22 M KOH?
a) 0.371 M
b) 0.455 M
c) 0.617 M
d) 0.742 M
129. What is the molarity of a Ba(OH)2 solution if 93.9 mL is completely titrated by 15.3 mL of
0.247 M H2SO4?
a) 0.0101 M
b) 0.0201 M
c) 0.0402 M
d) 0.0805 M
130. How is a Celsius temperature reading converted to a Kelvin temperature reading?
a) by adding 273
b) by subtracting 273
c) by dividing by 273
d) by multiplying by 273
131. The pH of a solution is 9. What is its H3O+ concentration?
a) 10-9 M
b) 10-7 M
-5
c) 10 M
d) 9 M
132. After balancing the equation FeCl3 + Zn  ZnCl2 + Fe, the coefficients, in order
from left to right, are
a) 2, 2, 1, 2.
b) 1, 1, 1, 1.
c) 4, 3, 3, 4.
d) 2, 3, 3, 2.
133. Which of the following is the electron configuration of carbon in the ground state?
a) 1s2 2s2 2p2
b) 2s2 2sl 2p3
2
2
c) 1s 2s 2p3
d) 1s2 2s2 2p6
134. How many valence electrons does a carbon atom have?
a) 3
b) 4
c) 5
d) 6
135. To draw the Lewis dot structure for Nitrogen, you would need to draw ___dots.
a) 7
b) 5
c) 3
d) 2
136. Which of the following would not speed up the rate that a reaction occurs?
a) catalyst
b) heat it up
c) use larger pieces
d) increase concentration
137. The shape of an ammonia (NH3) molecule is
a) trigonal pyramidal
b) bent
c) trigonal planar
d) tetrahedral
12