Final Review Part Two
Multiple Choice
Identify the choice that best completes the statement or answers the question.
____
1. Which hypothesis led to the discovery of the proton?
a. When a neutral hydrogen atom loses an electron, a positively-charged particle should
remain.
b. A proton should be 1840 times heavier than an electron.
c. Cathode rays should be attracted to a positively-charged plate.
d. The nucleus of an atom should contain neutrons.
____
2. Which of the following is correct concerning subatomic particles?
a. The electron was discovered by Goldstein in 1886.
b. The neutron was discovered by Chadwick in 1932.
c. The proton was discovered by Thomson in 1880.
d. Cathode rays were found to be made of protons.
____
3. All atoms are ____.
a. positively charged, with the number of protons exceeding the number of electrons
b. negatively charged, with the number of electrons exceeding the number of protons
c. neutral, with the number of protons equaling the number of electrons
d. neutral, with the number of protons equaling the number of electrons, which is equal to the
number of neutrons
____
4. Which of the following sets of symbols represents isotopes of the same element?
a.
c.
J
J
J
M
M
M
b.
d.
L
L
L
Q
Q
Q
____
5. How do the isotopes hydrogen-1 and hydrogen-2 differ?
a. Hydrogen-2 has one more electron than hydrogen-1.
b. Hydrogen-2 has one neutron; hydrogen-1 has none.
c. Hydrogen-2 has two protons; hydrogen-1 has one.
d. Hydrogen-2 has one proton; hydrogen-1 has none.
____
6. Which of the following isotopes has the same number of neutrons as phosphorus-31?
a.
c.
P
Si
b.
d.
S
Si
____
7. What is the maximum number of electrons in the second principal energy level?
a. 2
c. 18
b. 8
d. 32
____
8. When an electron moves from a lower to a higher energy level, the electron ____.
a. always doubles its energy
b. absorbs a continuously variable amount of energy
c. absorbs a quantum of energy
d. moves closer to the nucleus
____
9. If three electrons are available to fill three empty 2p atomic orbitals, how will the electrons be distributed in
the three orbitals?
a. one electron in each orbital
b. two electrons in one orbital, one in another, none in the third
c. three in one orbital, none in the other two
d. Three electrons cannot fill three empty 2p atomic orbitals.
____ 10. How many unpaired electrons are in a sulfur atom (atomic number 16)?
a. 0
c. 2
b. 1
d. 3
____ 11. How many half-filled orbitals are in a bromine atom?
a. 1
c. 3
b. 2
d. 4
____ 12. Which of the following electron configurations of outer sublevels is the most stable?
a. 4d 5s
c. 4d 5s
b. 4d 5s
d. 4d 5s
____ 13. What is the wavelength of an electromagnetic wave that travels at 3
MHz? (1 MHz = 1,000,000 Hz)
a.
10 m/s and has a frequency of 60
b. 60 MHz  300,000,000 m/s
c.
d. No answer can be determined from the information given.
____ 14. Which variable is directly proportional to frequency?
a. wavelength
c. position
b. velocity
d. energy
____ 15. How do the energy differences between the higher energy levels of an atom compare with the energy
differences between the lower energy levels of the atom?
a. They are greater in magnitude than those between lower energy levels.
b. They are smaller in magnitude than those between lower energy levels.
c. There is no significant difference in the magnitudes of these differences.
d. No answer can be determined from the information given.
____ 16. In an s orbital, the probability of finding an electron a particular distance from the nucleus does NOT depend
on ____.
a. a quantum mechanical model
c. the Schrodinger equation
b. direction with respect to the nucleus
d. the electron energy sublevel
____ 17. To what category of elements does an element belong if it is a poor conductor of electricity?
a. transition elements
c. nonmetals
b. metalloids
d. metals
____ 18. Of the elements Fe, Hg, U, and Te, which is a representative element?
a. Fe
c. U
b. Hg
d. Te
____ 19. Which of the following factors contributes to the increase in atomic size within a group in the periodic table
as the atomic number increases?
a. more shielding of the electrons by the highest occupied energy level
b. an increase in size of the nucleus
c. an increase in number of protons
d. fewer electrons in the highest occupied energy level
____ 20. Which of the following elements has the smallest atomic radius?
a. sulfur
c. selenium
b. chlorine
d. bromine
____ 21. In which of the following sets are the charges given correctly for all the ions?
a. Na , Mg , Al
c. Rb , Ba , P
b. K , Sr , O
d. N , O , F
____ 22. In which of the following groups of ions are the charges all shown correctly?
a. Li , O , S
c. K , F , Mg
b. Ca , Al , Br
d. Na , I , Rb
____ 23. Which of the following factors contributes to the increase in ionization energy from left to right across a
period?
a. an increase in the shielding effect
b. an increase in the size of the nucleus
c. an increase in the number of protons
d. fewer electrons in the highest occupied energy level
____ 24. As you move from left to right across the second period of the periodic table ____.
a. ionization energy increases
c. electronegativity decreases
b. atomic radii increase
d. atomic mass decreases
____ 25. Of the following elements, which one has the smallest first ionization energy?
a. boron
c. aluminum
b. carbon
d. silicon
____ 26. Which of the following pairs of elements is most likely to form an ionic compound?
a. magnesium and fluorine
b. nitrogen and sulfur
c. oxygen and chlorine
d. sodium and aluminum
____ 27. Which elements can form diatomic molecules joined by a single covalent bond?
a. hydrogen only
b. halogens only
c. halogens and members of the oxygen group only
d. hydrogen and the halogens only
____ 28. Which of the following electron configurations gives the correct arrangement of the four valence electrons of
the carbon atom in the molecule methane (CH )?
a. 2s 2p
c. 2s 2p 3s
b. 2s 2p 3s
d. 2s 2p
____ 29. Which of the following covalent bonds is the most polar?
a. H—F
b. H—C
c. H—H
d. H—N
____ 30. When placed between oppositely charged metal plates, the region of a water molecule attracted to the
negative plate is the ____.
a. hydrogen region of the molecule
c. H—O—H plane of the molecule
b. geometric center of the molecule
d. oxygen region of the molecule
____ 31. Which polyatomic ion forms a neutral compound when combined with a group 1A monatomic ion in a 1:1
ratio?
a. ammonium
c. nitrate
b. carbonate
d. phosphate
____ 32. Consider a mystery compound having the formula M T
.
If the compound is not an acid, if it contains only
two elements, and if M is not a metal, which of the following is true about the compound?
a. It contains a polyatomic ion.
c. Its name ends in -ic.
b. Its name ends in -ite or -ate.
d. It is a binary molecular compound.
____ 33. What is the formula for hydrosulfuric acid?
a. H S
b. H SO
c. HSO
d. H S
____ 34. What is the correct formula for barium chlorate?
a. Ba(ClO)
c. Ba(ClO )
b. Ba(ClO )
d. BaCl
____ 35. What is the correct formula for calcium dihydrogen phosphate?
a. CaH PO
c. Ca(H PO )
b. Ca H PO
d. Ca(H HPO )
____ 36. What is the correct name for Sn (PO ) ?
a. tritin diphosphate
b. tin(II) phosphate
c. tin(III) phosphate
d. tin(IV) phosphate
____ 37. Butanol is composed of carbon, hydrogen, and oxygen. If 1.0 mol of butanol contains 6.0
hydrogen, what is the subscript for the hydrogen atom in C H O?
a. 1
c. 6
b. 10
d. 8
10
atoms of
____ 38. Which of the following gas samples would have the largest number of representative particles at STP?
a. 12.0 L He
c. 0.10 L Xe
b. 7.0 L O
d. 0.007 L SO
____ 39. Given 1.00 mole of each of the following gases at STP, which gas would have the greatest volume?
a. He
c. SO
b. O
d. All would have the same volume.
____ 40. If the density of a noble gas is 1.783 g/L at STP, that gas is ____.
a. Kr
c. Ar
b. Xe
d. He
____ 41. Which of the following compounds has the lowest percent gold content by weight?
a. AuOH
c. AuCl
b. Au(OH)
d. AuI
____ 42. Which of the following compounds has the highest oxygen content, by weight?
a. Na O
c. BaO
b. CO
d. H O
____ 43. The ratio of carbon atoms to hydrogen atoms to oxygen atoms in a molecule of dicyclohexyl maleate is 4 to 6
to 1. What is its molecular formula if its molar mass is 280 g?
a. C H O
c. C H O
b. C H O
d. C H O
____ 44. Who was the man who lived from 460B.C.–370B.C. and was among the first to suggest the idea of atoms?
a. Atomos
c. Democritus
b. Dalton
d. Thomson
____ 45. Which of the following was NOT among Democritus’s ideas?
a. Matter consists of tiny particles called atoms.
b. Atoms are indivisible.
c. Atoms retain their identity in a chemical reaction.
d. Atoms are indestructible.
____ 46. Dalton's atomic theory included which idea?
a. All atoms of all elements are the same size.
b. Atoms of different elements always combine in one-to-one ratios.
c. Atoms of the same element are always identical.
d. Individual atoms can be seen with a microscope.
____ 47. Which of the following is NOT a part of Dalton's atomic theory?
a. All elements are composed of atoms.
b. Atoms are always in motion.
c. Atoms of the same element are identical.
d. Atoms that combine do so in simple whole-number ratios.
____ 48. Which of the following was originally a tenet of Dalton's atomic theory, but had to be revised about a century
ago?
a. Atoms are tiny indivisible particles.
b. Atoms of the same element are identical.
c. Compounds are made by combining atoms.
d. Atoms of different elements can combine with one another in simple whole number ratios.
____ 49. The comparison of the number of atoms in a copper coin the size of a penny with the number of people on
Earth is made to illustrate which of the following?
a. that atoms are indivisible
b. that atoms are very small
c. that atoms are very large
d. that in a copper penny, there is one atom for every person on Earth
____ 50. The range in size of most atomic radii is approximately ____.
a. 2 to 5 cm
c. 5 10
m to 2
10
m
b. 2 to 5 nm
d. 5
10
m to 2
10
m
____ 51. Why did J. J. Thomson reason that electrons must be a part of the atoms of all elements?
a. Cathode rays are negatively-charged particles.
b. Cathode rays can be deflected by magnets.
c. An electron is 2000 times lighter than a hydrogen atom.
d. Charge-to-mass ratio of electrons was the same, regardless of the gas used.
____ 52. Which of the following is true about subatomic particles?
a. Electrons are negatively charged and are the heaviest subatomic particle.
b. Protons are positively charged and the lightest subatomic particle.
c. Neutrons have no charge and are the lightest subatomic particle.
d. The mass of a neutron nearly equals the mass of a proton.
____ 53. Who conducted experiments to determine the quantity of charge carried by an electron?
a. Rutherford
c. Dalton
b. Millikan
d. Thomson
____ 54. What is the relative mass of an electron?
a. 1/1840 the mass of a hydrogen atom
b. 1/1840 the mass of a neutron + proton
c. 1/1840 the mass of a C-12 atom
d. 1/1840 the mass of an alpha particle
____ 55. The particles that are found in the nucleus of an atom are ____.
a. neutrons and electrons
c. protons and neutrons
b. electrons only
d. protons and electrons
____ 56. As a consequence of the discovery of the nucleus by Rutherford, which model of the atom is thought to be
true?
a. Protons, electrons, and neutrons are evenly distributed throughout the volume of the atom.
b. The nucleus is made of protons, electrons, and neutrons.
c. Electrons are distributed around the nucleus and occupy almost all the volume of the atom.
d. The nucleus is made of electrons and protons.
____ 57. The nucleus of an atom is ____.
a. the central core and is composed of protons and neutrons
b. positively charged and has more protons than neutrons
c. negatively charged and has a high density
d. negatively charged and has a low density
____ 58. In which of the following sets is the symbol of the element, the number of protons, and the number of
electrons given correctly?
a. In, 49 protons, 49 electrons
c. Cs, 55 protons, 132.9 electrons
b. Zn, 30 protons, 60 electrons
d. F, 19 protons, 19 electrons
____ 59. The mass number of an element is equal to ____.
a. the total number of electrons in the nucleus
b. the total number of protons and neutrons in the nucleus
c. less than twice the atomic number
d. a constant number for the lighter elements
____ 60. Using the periodic table, determine the number of neutrons in
a. 4
c. 16
b. 8
d. 24
O.
____ 61. How many protons, electrons, and neutrons does an atom with atomic number 50 and mass number 125
contain?
a. 50 protons, 50 electrons, 75 neutrons
c. 120 neutrons, 50 protons, 75 electrons
b. 75 electrons, 50 protons, 50 neutrons
d. 70 neutrons, 75 protons, 50 electrons
____ 62. Which of the following statements is NOT true?
a. Atoms of the same element can have different masses.
b. Atoms of isotopes of an element have different numbers of protons.
c. The nucleus of an atom has a positive charge.
d. Atoms are mostly empty space.
____ 63. If E is the symbol for an element, which two of the following symbols represent isotopes of the same
element?
1. E
2. E
3. E
4. E
a. 1 and 2
b. 3 and 4
c. 1 and 4
d. 2 and 3
____ 64. Select the correct symbol for an atom of tritium.
a.
c.
n
b.
d.
H
H
H
____ 65. How is the number of neutrons in the nucleus of an atom calculated?
a. Add the number of electrons and protons together.
b. Subtract the number of electrons from the number of protons.
c. Subtract the number of protons from the mass number.
d. Add the mass number to the number of electrons.
____ 66. In which of the following is the number of neutrons correctly represented?
a.
c.
F has 0 neutrons.
Mg has 24 neutrons.
b.
As has 108 neutrons.
d.
U has 146 neutrons.
____ 67. Which of the following statements is NOT true?
a. Protons have a positive charge.
b. Electrons are negatively charged and have a mass of 1 amu.
c. The nucleus of an atom is positively charged.
d. Neutrons are located in the nucleus of an atom.
____ 68. Why do chemists use relative masses of atoms compared to a reference isotope rather than the actual masses
of the atoms?
a. The actual mass of an electron is very large compared to the actual mass of a proton.
b. The actual masses of atoms are very small and difficult to work with.
c. The number of subatomic particles in atoms of different elements varies.
d. The actual masses of protons, electrons, and neutrons are not known.
____ 69. The atomic mass of an element is the ____.
a. total number of subatomic particles in its nucleus
b. weighted average of the masses of the isotopes of the element
c. total mass of the isotopes of the element
d. average of the mass number and the atomic number for the element
____ 70. The atomic mass of an element depends upon the ____.
a. mass of each electron in that element
b. mass of each isotope of that element
c. relative abundance of protons in that element
d. mass and relative abundance of each isotope of that element
____ 71. Which of the following is necessary to calculate the atomic mass of an element?
a. the atomic mass of carbon-12
b. the atomic number of the element
c. the relative masses of the element’s protons and neutrons
d. the masses of each isotope of the element
____ 72. The principal quantum number indicates what property of an electron?
a. position
c. energy level
b. speed
d. electron cloud shape
____ 73. Each period in the periodic table corresponds to ____.
a. a principal energy level
c. an orbital
b. an energy sublevel
d. a suborbital
____ 74. The modern periodic table is arranged in order of increasing atomic ____.
a. mass
c. number
b. charge
d. radius
____ 75. Of the elements Pt, V, Li, and Kr, which is a nonmetal?
a. Pt
c. Li
b. V
d. Kr
____ 76. The atomic number of an element is the total number of which particles in the nucleus?
a. neutrons
c. electrons
b. protons
d. protons and electrons
____ 77. What element has the electron configuration 1s 2s 2p 3s 3p ?
a. nitrogen
c. silicon
b. selenium
d. silver
____ 78. Which of the following is true about the electron configurations of the noble gases?
a. The highest occupied s and p sublevels are completely filled.
b. The highest occupied s and p sublevels are partially filled.
c. The electrons with the highest energy are in a d sublevel.
d. The electrons with the highest energy are in an f sublevel.
____ 79. Elements that are characterized by the filling of p orbitals are classified as ____.
a. groups 3A through 8A
c. inner transition metals
b. transition metals
d. groups 1A and 2A
____ 80. Which of the following electron configurations is most likely to result in an element that is relatively
inactive?
a. a half-filled energy sublevel
b. a filled energy sublevel
c. one empty and one filled energy sublevel
d. a filled highest occupied principal energy level
____ 81. Which subatomic particle plays the greatest part in determining the properties of an element?
a. proton
c. neutron
b. electron
d. none of the above
____ 82. Which of the following is true about the electron configurations of the representative elements?
a. The highest occupied s and p sublevels are completely filled.
b. The highest occupied s and p sublevels are partially filled.
c. The electrons with the highest energy are in a d sublevel.
d. The electrons with the highest energy are in an f sublevel.
____ 83. What are the Group 1A and Group 7A elements examples of?
a. representative elements
c. noble gases
b. transition elements
d. nonmetallic elements
____ 84. How does atomic radius change from left to right across a period in the periodic table?
a. It tends to decrease.
c. It first increases, then decreases.
b. It tends to increase.
d. It first decreases, then increases.
____ 85. What causes the shielding effect to remain constant across a period?
a. Electrons are added to the same principal energy level.
b. Electrons are added to different principal energy levels.
c. The charge on the nucleus is constant.
d. The atomic radius increases.
____ 86. Atomic size generally ____.
a. increases as you move from left to right across a period
b. decreases as you move from top to bottom within a group
c. remains constant within a period
d. decreases as you move from left to right across a period
____ 87. What element in the second period has the largest atomic radius?
a. carbon
c. potassium
b. lithium
d. neon
____ 88. Which of the following statements is true about ions?
a. Cations form when an atom gains electrons.
b. Cations form when an atom loses electrons.
c. Anions form when an atom gains protons.
d. Anions form when an atom loses protons.
____ 89. The metals in Groups 1A, 2A, and 3A ____.
a. gain electrons when they form ions
b. all form ions with a negative charge
c. all have ions with a 1 charge
d. lose electrons when they form ions
____ 90. Which of the following statements is NOT true about ions?
a. Cations are positively charged ions.
b. Anions are common among nonmetals.
c. Charges for ions are written as numbers followed by a plus or minus sign.
d. When a cation forms, more electrons are transferred to it.
____ 91. Why is the second ionization energy greater than the first ionization energy?
a. It is more difficult to remove a second electron from an atom.
b. The size of atoms increases down a group.
c. The size of anions decreases across a period.
d. The nuclear attraction from protons in the nucleus decreases.
____ 92. Which of the following elements has the smallest ionic radius?
a. Li
c. O
b. K
d. S
____ 93. What is the energy required to remove an electron from an atom in the gaseous state called?
a. nuclear energy
c. shielding energy
b. ionization energy
d. electronegative energy
____ 94. For Group 2A metals, which electron is the most difficult to remove?
a. the first
b. the second
c. the third
d. All the electrons are equally difficult to remove.
____ 95. Which of the following factors contributes to the decrease in ionization energy within a group in the periodic
table as the atomic number increases?
a. increase in atomic size
b. increase in size of the nucleus
c. increase in number of protons
d. fewer electrons in the highest occupied energy level
____ 96. Which of the following elements has the smallest first ionization energy?
a. sodium
c. potassium
b. calcium
d. magnesium
____ 97. Which of the following elements has the lowest electronegativity?
a. lithium
c. bromine
b. carbon
d. fluorine
____ 98. Which statement is true about electronegativity?
a. Electronegativity is the ability of an anion to attract another anion.
b. Electronegativity generally increases as you move from top to bottom within a group.
c. Electronegativity generally is higher for metals than for nonmetals.
d. Electronegativity generally increases from left to right across a period.
____ 99. Compared with the electronegativities of the elements on the left side of a period, the electronegativities of
the elements on the right side of the same period tend to be ____.
a. lower
c. the same
b. higher
d. unpredictable
____ 100. Which of the following decreases with increasing atomic number in Group 2A?
a. shielding effect
c. ionization energy
b. ionic size
d. number of electrons
____ 101. Which of the following statements correctly compares the relative size of an ion to its neutral atom?
a. The radius of an anion is greater than the radius of its neutral atom.
b. The radius of an anion is identical to the radius of its neutral atom.
c. The radius of a cation is greater than the radius of its neutral atom.
d. The radius of a cation is identical to the radius of its neutral atom.
____ 102. What is the electron configuration of the calcium ion?
a. 1s 2s 2p 3s 3p
b. 1s 2s 2p 3s 3p 4s
c. 1s 2s 2p 3s 3p 4s
d. 1s 2s 2p 3s
____ 103. What is the electron configuration of the gallium ion?
a. 1s 2s 2p 3s 3p
c. 1s 2s 2p 3s 3p 4s 4p
b. 1s 2s 2p 3s 3p 4s
d. 1s 2s 2p 3s 3p 3d
____ 104. The octet rule states that, in chemical compounds, atoms tend to have ____.
a. the electron configuration of a noble gas
b. more protons than electrons
c. eight electrons in their principal energy level
d. more electrons than protons
____ 105. What is the formula of the ion formed when tin achieves a stable electron configuration?
a. Sn
c. Sn
b. Sn
d. Sn
____ 106. What is the formula of the ion formed when cadmium achieves a pseudo-noble-gas electron configuration?
a. Cd
c. Cd
b. Cd
d. Cd
____ 107. Which of the following is a pseudo-noble-gas electron configuration?
a. 1s 2s 2p 3s 3d
c. 1s 2s 2p 3s 3p 3d
b. 1s 2s 2p 3s 3p
d. 1s 2s 2p 3s 3d 4s
____ 108. What is the electron configuration of the oxide ion (O )?
a. 1s 2s 2p
c. 1s 2s
b. 1s 2s 2p
d. 1s 2s 2p
____ 109. What is the electron configuration of the iodide ion?
a. 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s 5p
b. 1s 2s 2p 3s 3p 3d 4s 4p 4d
c. 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s
d. 1s 2s 2p 3s 3p 3d 4s 4p
____ 110. How many valence electrons are transferred from the nitrogen atom to potassium in the formation of the
compound potassium nitride?
a. 0
c. 2
b. 1
d. 3
____ 111. How many valence electrons are transferred from the calcium atom to iodine in the formation of the
compound calcium iodide?
a. 0
c. 2
b. 1
d. 3
____ 112. What is the formula unit of sodium nitride?
a. NaN
b. Na N
c. Na N
d. NaN
____ 113. What is the formula unit of aluminum oxide?
a. AlO
b. Al O
c. AlO
d. Al O
____ 114. What is the name of the ionic compound formed from lithium and bromine?
a. lithium bromine
c. lithium bromium
b. lithium bromide
d. lithium bromate
____ 115. What is the formula for sodium sulfate?
a. NaSO
b. Na SO
c. Na(SO )
d. Na (SO )
____ 116. Alloys are commonly used in manufacturing. Which of the following is NOT a reason to use an alloy instead
of a pure metal?
a. Bronze is tougher than pure copper.
c. Brass is more malleable than pure copper.
b. Sterling silver is stronger than pure silver. d. Cast iron is more brittle than pure iron.
____ 117. Which of the following compounds has the formula KNO ?
a. potassium nitrate
c. potassium nitrite
b. potassium nitride
d. potassium nitrogen oxide
____ 118. Which is a typical characteristic of an ionic compound?
a. Electron pairs are shared among atoms.
b. The ionic compound has a low solubility in water.
c. The ionic compound is described as a molecule.
d. The ionic compound has a high melting point.
____ 119. How do atoms achieve noble-gas electron configurations in single covalent bonds?
a. One atom completely loses two electrons to the other atom in the bond.
b. Two atoms share two pairs of electrons.
c. Two atoms share two electrons.
d. Two atoms share one electron.
____ 120. Why do atoms share electrons in covalent bonds?
a. to become ions and attract each other
b. to attain a noble-gas electron configuration
c. to become more polar
d. to increase their atomic numbers
____ 121. Which of the following elements can form diatomic molecules held together by triple covalent bonds?
a. carbon
c. fluorine
b. oxygen
d. nitrogen
____ 122. Which noble gas has the same electron configuration as the oxygen in a water molecule?
a. helium
c. argon
b. neon
d. xenon
____ 123. Which of the following diatomic molecules is joined by a double covalent bond?
a.
c.
b.
d.
____ 124. A molecule with a single covalent bond is ____.
a. CO
b. Cl
c. CO
d. N
____ 125. When one atom contributes both bonding electrons in a single covalent bond, the bond is called a(n) ____.
a. one-sided covalent bond
c. coordinate covalent bond
b. unequal covalent bond
d. ionic covalent bond
____ 126. Once formed, how are coordinate covalent bonds different from other covalent bonds?
a. They are stronger.
c. They are weaker.
b. They are more ionic in character.
d. There is no difference.
____ 127. When H forms a bond with H O to form the hydronium ion H O , this bond is called a coordinate
covalent bond because ____.
a. both bonding electrons come from the oxygen atom
b. it forms an especially strong bond
c. the electrons are equally shared
d. the oxygen no longer has eight valence electrons
____ 128. Which of the following bonds is the least reactive?
a. C—C
c. O—H
b. H—H
d. H—Cl
____ 129. In which of the following compounds is the octet expanded to include 12 electrons?
a. H S
c. PCl
b. PCl
d. SF
____ 130. How is a pair of molecular orbitals formed?
a. by the splitting of a single atomic orbital
b. by the reproduction of a single atomic orbital
c. by the overlap of two atomic orbitals from the same atom
d. by the overlap of two atomic orbitals from different atoms
____ 131. The side-by-side overlap of p orbitals produces what kind of bond?
a. alpha bond
c. pi bond
b. beta bond
d. sigma bond
____ 132. Where are the electrons most probably located in a molecular bonding orbital?
a. anywhere in the orbital
b. between the two atomic nuclei
c. in stationary positions between the two atomic nuclei
d. in circular orbits around each nucleus
____ 133. Sigma bonds are formed as a result of the overlapping of which type(s) of atomic orbital(s)?
a. s only
c. d only
b. p only
d. s and p
____ 134. Which of the following bond types is normally the weakest?
a. sigma bond formed by the overlap of two s orbitals
b. sigma bond formed by the overlap of two p orbitals
c. sigma bond formed by the overlap of one s and one p orbital
d. pi bond formed by the overlap of two p orbitals
____ 135. What causes water molecules to have a bent shape, according to VSEPR theory?
a. repulsive forces between unshared pairs of electrons
b. interaction between the fixed orbitals of the unshared pairs of oxygen
c. ionic attraction and repulsion
d. the unusual location of the free electrons
____ 136. What type of hybrid orbital exists in the methane molecule?
a. sp
c. sp
b. sp
d. sp d
____ 137. What is the shape of a molecule with a triple bond?
a. tetrahedral
c. bent
b. pyramidal
d. linear
____ 138. What type of hybridization occurs in the orbitals of a carbon atom participating in a triple bond with another
carbon atom?
a.
c.
b.
d.
____ 139. How many pi bonds are formed when sp hybridization occurs in ethene, C H ?
a. 0
c. 2
b. 1
d. 3
____ 140. Which of the following atoms acquires the most negative charge in a covalent bond with hydrogen?
a. C
c. O
b. Na
d. S
____ 141. A bond formed between a silicon atom and an oxygen atom is likely to be ____.
a. ionic
c. polar covalent
b. coordinate covalent
d. nonpolar covalent
____ 142. What causes hydrogen bonding?
a. attraction between ions
b. motion of electrons
c. sharing of electron pairs
d. bonding of a covalently bonded hydrogen atom with an unshared electron pair
____ 143. Why is hydrogen bonding only possible with hydrogen?
a. Hydrogen’s nucleus is electron deficient when it bonds with an electronegative atom.
b. Hydrogen is the only atom that is the same size as an oxygen atom.
c. Hydrogen is the most electronegative element.
d. Hydrogen tends to form covalent bonds.
____ 144. In which of the following are the symbol and name for the ion given correctly?
a. Fe : ferrous ion; Fe : ferric ion
b. Sn : stannic ion; Sn : stannous ion
c. Co : cobalt(II) ion; Co : cobaltous ion
d. Pb : lead ion; Pb : lead(IV) ion
____ 145. Which of the following correctly provides the name of the element, the symbol for the ion, and the name of
the ion?
a. fluorine, F , fluoride ion
b. zinc, Zn , zincate ion
c. copper, Cu , cuprous ion
d. sulfur, S , sulfurous ion
____ 146. The nonmetals in Groups 6A and 7A ____.
a. lose electrons when they form ions
b. have a numerical charge that is found by subtracting 8 from the group number
c. all have ions with a –1 charge
d. end in -ate
____ 147. What determines that an element is a metal?
a. the magnitude of its charge
b. the molecules that it forms
c. when it is a Group A element
d. its position in the periodic table
____ 148. What is the Stock name for chromic ion?
a. chromium(I) ion
b. chromium(II) ion
c. chromium(III) ion
d. chromium(IV) ion
____ 149. In which of the following are the symbol and name for the ion given correctly?
a. NH : ammonia; H : hydride
c. OH : hydroxide; O : oxide
b. C H O
: acetate; C O
: oxalite
d. PO
: phosphate; PO
: phosphite
____ 150. Which of the following correctly provides the names and formulas of polyatomic ions?
a. carbonate: HCO ; bicarbonate: CO
b. nitrite: NO ; nitrate: NO
c. sulfite: S ; sulfate: SO
d. chromate: CrO
; dichromate: Cr O
____ 151. An -ate or -ite at the end of a compound name usually indicates that the compound contains ____.
a. fewer electrons than protons
c. only two elements
b. neutral molecules
d. a polyatomic anion
____ 152. Which of the following is true about the composition of ionic compounds?
a. They are composed of anions and cations.
b. They are composed of anions only.
c. They are composed of cations only.
d. They are formed from two or more nonmetallic elements.
____ 153. Which of the following formulas represents an ionic compound?
a. CS
c. N O
b. BaI
d. PCl
____ 154. Which element, when combined with fluorine, would most likely form an ionic compound?
a. lithium
c. phosphorus
b. carbon
d. chlorine
____ 155. Which of the following shows correctly an ion pair and the ionic compound the two ions form?
a. Sn , N ; Sn N
c. Cr , I ; CrI
b. Cu , O ; Cu O
d. Fe , O ; Fe O
____ 156. In which of the following are the formula of the ionic compound and the charge on the metal ion shown
correctly?
a. UCl , U
c. IrS , Ir
b. ThO , Th
d. NiO, Ni
____ 157. Which of the following correctly represents an ion pair and the ionic compound the ions form?
a. Ca , F ; CaF
c. Ba , O ; Ba O
b. Na , Cl ; NaCl
d. Pb , O ; Pb O
____ 158. In which of the following is the name and formula given correctly?
a. sodium oxide, NaO
c. cobaltous chloride, CoCl
b. barium nitride, BaN
d. stannic fluoride, SnF
____ 159. Which of the following compounds contains the lead(II) ion?
a. PbO
c. Pb2O
b. PbCl4
d. Pb2S
____ 160. Which set of chemical name and chemical formula for the same compound is correct?
a. iron(II) oxide, Fe O
c. tin(IV) bromide, SnBr
b. aluminum fluorate, AlF
d. potassium chloride, K Cl
____ 161. What is the correct formula for potassium sulfite?
a. KHSO
c. K SO
b. KHSO
d. K SO
____ 162. Which set of chemical name and chemical formula for the same compound is correct?
a. ammonium sulfite, (NH ) S
c. lithium carbonate, LiCO
b. iron(III) phosphate, FePO
d. magnesium dichromate, MgCrO
____ 163. What type of compound is CuSO ?
a. monotomic ionic
b. polyatomic covalent
c. polyatomic ionic
d. binary molecular
____ 164. Sulfur hexafluoride is an example of a ____.
a. monatomic ion
b. polyatomic ion
c. binary compound
d. polyatomic compound
____ 165. Which of the following correctly shows a prefix used in naming binary molecular compounds with its
corresponding number?
a. deca-, 7
c. hexa-, 8
b. nona-, 9
d. octa-, 4
____ 166. Which of the following is a binary molecular compound?
a. BeHCO
c. AgI
b. PCl
d. MgS
____ 167. Which of the following formulas represents a molecular compound?
a. ZnO
c. SO
b. Xe
d. BeF
____ 168. When naming acids, the prefix hydro- is used when the name of the acid anion ends in ____.
a. -ide
c. -ate
b. -ite
d. -ic
____ 169. Which of the following shows both the correct formula and correct name of an acid?
a. HClO , chloric acid
c. H PO , phosphoric acid
b. HNO , hydronitrous acid
d. HI, iodic acid
____ 170. What is the name of H SO ?
a. hyposulfuric acid
b. hydrosulfuric acid
c. sulfuric acid
d. sulfurous acid
____ 171. When the name of an anion that is part of an acid ends in -ite, the acid name includes the suffix ____.
a. -ous
c. -ate
b. -ic
d. -ite
____ 172. What is the formula for sulfurous acid?
a. H SO
b. H SO
c. H SO
d. H S
____ 173. What is the formula for phosphoric acid?
a. H PO
b. H PO
c. HPO
d. HPO
____ 174. Which of the following pairs of substances best illustrates the law of multiple proportions?
a. H and O
c. CaCl and CaBr
b. P O and PH
d. NO and NO
____ 175. Select the correct formula for sulfur hexafluoride.
a. S F
c. F S
b. F SO
d. SF
____ 176. What is the correct name for the compound CoCl ?
a. cobalt(I) chlorate
c. cobalt(II) chlorate
b. cobalt(I) chloride
d. cobalt(II) chloride
____ 177. Suppose you encounter a chemical formula with H as the cation. What do you know about this compound
immediately?
a. It is a polyatomic ionic compound.
c. It is a base.
b. It is an acid.
d. It has a 1 charge.
____ 178. Which of the following is the correct name for N O ?
a. nitrous oxide
c. nitrogen dioxide
b. dinitrogen pentoxide
d. nitrate oxide
____ 179. How many moles of tungsten atoms are in 4.8 10 atoms of tungsten?
a. 8.0 10 moles
c. 1.3 10 moles
b. 8.0 10 moles
d. 1.3 10 moles
____ 180. How many moles of silver atoms are in 1.8
10
atoms of silver?
a. 3.0
b. 3.3
c. 3.0
d. 1.1
10
10
____ 181. How many atoms are in 0.075 mol of titanium?
a. 1.2 10-25
c. 6.4
b. 2.2 10
d. 4.5
10
10
10
10
____ 182. How many molecules are in 2.10 mol CO ?
a. 2.53 10 molecules
b. 3.79 10 molecules
c. 3.49
d. 1.26
____ 183. How many atoms are in 3.5 moles of arsenic atoms?
a. 5.8 10
c. 2.1
atoms
b. 7.5 10 atoms
d. 1.7
10
10
10
10
molecules
molecules
atoms
atoms
____ 184. Which of the following is NOT a true about atomic mass?
a. The atomic mass is 12 g for magnesium.
b. The atomic mass is the mass of one mole of atoms.
c. The atomic mass is found by checking the periodic table.
d. The atomic mass is the number of grams of an element that is numerically equal to the
mass in amu.
____ 185. What is true about the molar mass of chlorine gas?
a. The molar mass is 35.5 g.
b. The molar mass is 71.0 g.
c. The molar mass is equal to the mass of one mole of chlorine atoms.
d. none of the above
____ 186. What is the molar mass of AuCl3?
a. 96 g
b. 130 g
c. 232.5 g
d. 303.6 g
____ 187. What is the molar mass of (NH ) CO ?
a. 144 g
b. 138 g
c. 96 g
d. 78 g
____ 188. The molar mass of C H
a. carbon atoms
b. anions
and the molar mass of CaCO contain approximately the same number of ____.
c. cations
d. grams
____ 189. What is the mass in grams of 5.90 mol C H ?
a. 0.0512 g
c. 389 g
b. 19.4 g
d. 673 g
____ 190. What is the number of moles in 432 g Ba(NO ) ?
a. 0.237 mol
c. 1.65 mol
b. 0.605 mol
d. 3.66 mol
____ 191. What is the number of moles of beryllium atoms in 36 g of Be?
a. 0.25 mol
c. 45.0 mol
b. 4.0 mol
d. 320 mol
____ 192. How many moles of CaBr are in 5.0 grams of CaBr ?
a. 2.5 10 mol
c. 4.0 10 mol
b. 4.2 10 mol
d. 1.0 10 mol
____ 193. For which of the following conversions does the value of the conversion factor depend upon the formula of
the substance?
a. volume of gas (STP) to moles
b. density of gas (STP) to molar mass
c. mass of any substance to moles
d. moles of any substance to number of particles
____ 194. What is the mass of silver in 3.4 g AgNO ?
a. 0.025 g
b. 0.64 g
c. 2.2 g
d. 3.0 g
____ 195. What is the mass of oxygen in 250 g of sulfuric acid, H SO ?
a. 0.65 g
c. 16 g
b. 3.9 g
d. 160 g
____ 196. What is the volume, in liters, of 0.500 mol of C H gas at STP?
a. 0.0335 L
c. 16.8 L
b. 11.2 L
d. 22.4 L
____ 197. What is the number of moles in 500 L of He gas at STP?
a. 0.05 mol
c. 22 mol
b. 0.2 mol
d. 90 mol
____ 198. What is the number of moles in 9.63 L of H S gas at STP?
a. 0.104 mol
c. 3.54 mol
b. 0.430 mol
d. 14.7 mol
____ 199. What is the density at STP of the gas sulfur hexafluoride, SF ?
a. 0.153 g/L
c. 3270 g/L
b. 6.52 g/L
d. 3.93 10 g/L
____ 200. The molar mass of a certain gas is 49 g. What is the density of the gas in g/L at STP?
a. 3.6 10
c. 2.2 g/L
g/L
b. 0.46 g/L
d. 71 g/L
____ 201. A 22.4-L sample of which of the following substances, at STP, would contain 6.02
particles?
a. oxygen
c. cesium iodide
b. gold
d. sulfur
10
representative
____ 202. If the density of an unknown gas Z is 4.50 g/L at STP, what is the molar mass of gas Z?
a. 0.201 g/mol
c. 26.9 g/mol
b. 5.00 g/mol
d. 101 g/mol
____ 203. If 60.2 grams of Hg combines completely with 24.0 grams of Br to form a compound, what is the percent
composition of Hg in the compound?
a. 28.5%
c. 71.5%
b. 39.9%
d. 60.1%
____ 204. What is the percent composition of chromium in BaCrO ?
a. 4.87%
c. 20.5%
b. 9.47%
d. 25.2%
____ 205. If 20.0 grams of Ca combines completely with 16.0 grams of S to form a compound, what is the percent
composition of Ca in the compound?
a. 1.25%
c. 44.4%
b. 20.0%
d. 55.6%
____ 206. What is the percent composition of carbon, in heptane, C H ?
a. 12%
c. 68%
b. 19%
d. 84%
____ 207. What is the percent by mass of carbon in acetone, C H O?
a. 20.7%
c. 1.61%
b. 62.1%
d. 30.0%
____ 208. Which expression represents the percent by mass of nitrogen in NH4NO3?
a. 14 g N/80 g NH NO
c. 80 g NH NO /14 g N
100%
b. 28 g N/80 g NH NO
d. 80 g NH NO /28 g N
100%
100%
100%
____ 209. What is the empirical formula of a compound that is 40% sulfur and 60% oxygen by weight?
a. SO
c. SO
b. SO
d. S O
____ 210. What is the empirical formula of a substance that is 53.5% C, 15.5% H, and 31.1% N by weight?
a. C HN
c. C H N
b. C H N
d. CH N
____ 211. Which of the following sets of empirical formula, molar mass, and molecular formula is correct?
a. CH, 78 g, C H
c. CaO, 56 g, Ca O
b. CH N, 90 g, C H N
d. C H O, 120 g, C H O
____ 212. Which of the following is the correct skeleton equation for the reaction that takes place when solid
phosphorus combines with oxygen gas to form diphosphorus pentoxide?
a. P(s) O (g)  PO (g)
c. P(s) O2(g)  P O (s)
b. P(s) O(g)  P O (g)
d. P O (s)  P (s) O (g)
____ 213. What are the missing coefficients for the skeleton equation below?
Cr(s) Fe(NO ) (aq)
Fe(s) Cr(NO ) (aq)
a. 4, 6, 6, 2
c. 2, 3, 3, 2
b. 2, 3, 2, 3
d. 1, 3, 3, 1
____ 214. What are the missing coefficients for the skeleton equation below?
Al (SO ) (aq) KOH(aq)  Al(OH) (aq) K SO (aq)
a. 1, 3, 2, 3
c. 4, 6, 2, 3
b. 2, 12, 4, 6
d. 1, 6, 2, 3
____ 215. When potassium hydroxide and barium chloride react, potassium chloride and barium hydroxide are formed.
The balanced equation for this reaction is ____.
a. KH BaCl
KCl BaH
c. 2KOH BaCl
2KCl Ba(OH)
b. KOH BaCl
KCl BaOH
d. KOH BaCl
KCl
BaOH
____ 216. Which of the following is a balanced equation representing the decomposition of lead(IV) oxide?
a. PbO
c. Pb O
Pb 2O
2Pb O
b. PbO
d. PbO
Pb O
Pb O
____ 217. What are the correct formulas and coefficients for the products of the following double-replacement reaction?
RbOH H PO

a. Rb(PO )
c. Rb PO
H O
3H O
b. RbPO
d. H Rb PO OH
2H O
____ 218. When the equation for the complete combustion of one mole of C H OH is balanced, the coefficient for
oxygen is ____.
a. 13
c. 7
2
2
b. 11
d. 9
2
2
____ 219. Which of the following statements is NOT true about the decomposition of a simple binary compound?
a. The products are unpredictable.
b. The products are the constituent elements.
c. The reactant is a single substance.
d. The reactant could be an ionic or a molecular compound.
____ 220. Which of the following statements is true about single-replacement reactions?
a. They are restricted to metals.
c. Two reactants produce two products.
b. They involve a single product.
d. Any metal replaces any other metal.
____ 221. In the activity series of metals, which metal(s) will displace hydrogen from an acid?
a. only metals above hydrogen
c. any metal
b. only metals below hydrogen
d. only metals from Li to Na
____ 222. Use the activity series of metals to complete a balanced chemical equation for the following single
replacement reaction.
Ag(s) KNO (aq) 
a. AgNO
K
b. AgK NO
c. AgKNO
d. No reaction takes place because silver is less reactive than potassium.
____ 223. Which of the following statements is NOT true about double-replacement reactions?
a. The product may precipitate from solution.
b. The product may be a gas.
c. The product may be a molecular compound.
d. The reactant may be a solid metal.
____ 224. In a double-replacement reaction, ____.
a.
b.
c.
d.
the reactants are usually a metal and a nonmetal
one of the reactants is often water
the reactants are generally two ionic compounds in aqueous solution
energy in the form of heat or light is often produced
____ 225. A double-replacement reaction takes place when aqueous Na CO reacts with aqueous Sn(NO ) . You
would expect one of the products of this reaction to be ____.
a. NaNO
c. Sn(CO )
b. NaSn
d. CNO
____ 226. The complete combustion of which of the following substances produces carbon dioxide and water?
a. C H
c. CaHCO
b. K CO
d. NO
____ 227. Which of the following is the correctly balanced equation for the incomplete combustion of heptene, C H ?
a. C H
c. 2C H
14O
7CO 7H O
21O
14CO
14H O
b. C H
d. C H
7O
7CO 7H O
O
C O
7H
____ 228. If a combination reaction takes place between rubidium and bromine, the chemical formula for the product is
____.
a. RuBr
c. RbBr
b. Rb Br
d. RbBr
____ 229. What is the balanced chemical equation for the reaction that takes place between bromine and sodium iodide?
a. Br
c. Br NaI
NaI
NaBr
I
NaBrI
b. Br
d. Br NaI
2NaI
2NaBr I
NaBr I
____ 230. What is the driving force in the following reaction?
Ni(NO ) (aq) K S(aq)
NiSs 2KNO (aq)
a. A gas is formed.
c. Ionic compounds are reactants.
b. A precipitate is formed.
d. Ionic compounds are products.
Final Review Part Two
Answer Section
MULTIPLE CHOICE
1. ANS: A
PTS: 1
DIF: L3
REF: p. 106
OBJ: 4.2.1 Identify three types of subatomic particle.
STA: 5.2.B.1
2. ANS: B
PTS: 1
DIF: L3
REF: p. 105 | p. 106
OBJ: 4.2.1 Identify three types of subatomic particle.
STA: 5.2.B.1
3. ANS: C
PTS: 1
DIF: L3
REF: p. 106
OBJ: 4.2.1 Identify three types of subatomic particle.
STA: 5.6.A.1
4. ANS: C
PTS: 1
DIF: L3
REF: p. 112 | p. 113
OBJ: 4.3.1 Explain what makes elements and isotopes different from each other.
5. ANS: B
PTS: 1
DIF: L3
REF: p. 111 | p. 112 | p. 113
OBJ: 4.3.1 Explain what makes elements and isotopes different from each other. | 4.3.2 Calculate the
number of neutrons in an atom.
6. ANS: B
PTS: 1
DIF: L3
REF: p. 111
OBJ: 4.3.2 Calculate the number of neutrons in an atom.
7. ANS: B
PTS: 1
DIF: L3
REF: p. 132
OBJ: 5.1.3 Describe the energies and positions of electrons according to the quantum mechanical model.
STA: 5.6.A.8
8. ANS: C
PTS: 1
DIF: L3
REF: p. 128
OBJ: 5.1.3 Describe the energies and positions of electrons according to the quantum mechanical model.
STA: 5.6.A.8
9. ANS: A
PTS: 1
DIF: L3
REF: p. 134
OBJ: 5.2.1 Describe how to write the electron configuration for an atom.
STA: 5.6.A.3
10. ANS: C
PTS: 1
DIF: L3
REF: p. 133 | p. 134
OBJ: 5.2.1 Describe how to write the electron configuration for an atom.
STA: 5.6.A.3
11. ANS: A
PTS: 1
DIF: L3
REF: p. 133 | p. 134
OBJ: 5.2.1 Describe how to write the electron configuration for an atom.
STA: 5.6.A.3
12. ANS: A
PTS: 1
DIF: L3
REF: p. 133 | p. 134 | p. 135 | p. 136
OBJ: 5.2.2 Explain why the actual electron configurations for some elements differ from those predicted by
the aufbau principle.
STA: 5.6.A.8
13. ANS: C
PTS: 1
DIF: L3
REF: p. 140
OBJ: 5.3.1 Describe the relationship between the wavelength and frequency of light.
14. ANS: D
PTS: 1
DIF: L3
REF: p. 142
OBJ: 5.3.3 Explain how the frequencies of light are related to changes in electron energies.
15. ANS: B
PTS: 1
DIF: L3
REF: p. 143
OBJ: 5.3.3 Explain how the frequencies of light are related to changes in electron energies.
16. ANS: B
PTS: 1
DIF: L3
REF: p. 131
OBJ: 5.1.4 Describe how the shapes of orbitals at different sublevels differ. | 5.3.4 Distinguish between
quantum mechanics and classical mechanics.
STA: 5.6.A.8
17. ANS: C
PTS: 1
DIF: L3
REF: p. 160
OBJ: 6.1.3 Identify three broad classes of elements.
STA: 5.6.A.5
18. ANS: D
PTS: 1
DIF: L3
REF: p. 162 | p. 163 | p. 164
19.
20.
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
31.
32.
33.
34.
35.
OBJ: 6.2.3 Distinguish representative elements and transition metals.
STA: 5.6.A.5
ANS: A
PTS: 1
DIF: L3
REF: p. 171
OBJ: 6.3.1 Describe trends among elements for atomic size.
STA: 5.6.A.5
ANS: B
PTS: 1
DIF: L3
REF: p. 171 | p. 175
OBJ: 6.3.1 Describe trends among elements for atomic size.
STA: 5.6.A.5
ANS: B
PTS: 1
DIF: L3
REF: p. 162 | p. 163 | p. 172
OBJ: 6.3.2 Explain how ions form.
STA: 5.6.A.4
ANS: B
PTS: 1
DIF: L3
REF: p. 162 | p. 163 | p. 172
OBJ: 6.3.2 Explain how ions form.
STA: 5.6.A.4
ANS: C
PTS: 1
DIF: L3
REF: p. 174
OBJ: 6.3.3 Describe periodic trends for first ionization energy, ionic size, and electronegativity.
STA: 5.6.A.5 | 5.6.A.8
ANS: A
PTS: 1
DIF: L3
REF: p. 178
OBJ: 6.2.1 Describe the information in a periodic table. | 6.3.3 Describe periodic trends for first ionization
energy, ionic size, and electronegativity.
STA: 5.6.A.5
ANS: C
PTS: 1
DIF: L3
REF: p. 173
OBJ: 6.3.3 Describe periodic trends for first ionization energy, ionic size, and electronegativity.
STA: 5.6.A.5 | 5.6.A.8
ANS: A
PTS: 1
DIF: L3
REF: p. 194
OBJ: 7.2.1 Explain the electrical charge of an ionic compound.
STA: 5.6.A.3
ANS: D
PTS: 1
DIF: L3
REF: p. 217 | p. 218
OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet
rule. STA:
5.6.A.4
ANS: D
PTS: 1
DIF: L3
REF: p. 220 | p. 234
OBJ: 8.2.2 Demonstrate how electron dot structures represent shared electrons.
STA: 5.6.A.3
ANS: A
PTS: 1
DIF: L3
REF: p. 238 | p. 239
OBJ: 8.4.1 Describe how electronegativity values determine the charge distribution in a polar molecule.
STA: 5.6.A.4
ANS: A
PTS: 1
DIF: L3
REF: p. 239
OBJ: 8.4.2 Describe what happens to polar molecules when they are placed between oppositely charged
metal plates.
ANS: C
PTS: 1
DIF: L3
REF: p. 257 | p. 264
OBJ: 9.1.2 Define a polyatomic ion and write the names and formulas of the most common polyatomic
ions. | 9.2.2 Apply the rules for naming and writing formulas for compounds with polyatomic ions.
STA: 5.6.A.5
ANS: D
PTS: 1
DIF: L3
REF: p. 268 | p. 269
OBJ: 9.3.2 Apply the rules for naming and writing formulas for binary molecular compounds.
STA: 5.6.A.5
ANS: D
PTS: 1
DIF: L3
REF: p. 272
OBJ: 9.4.2 Apply the rules in reverse to write formulas of acids.
STA: 5.6.A.5
ANS: C
PTS: 1
DIF: L3
REF: p. 257 | p. 264
OBJ: 9.2.2 Apply the rules for naming and writing formulas for compounds with polyatomic ions. | 9.5.2
Apply the rules for naming chemical compounds by using a flowchart.
STA: 5.6.A.5
ANS: C
PTS: 1
DIF: L3
REF: p. 257 | p. 264
36.
37.
38.
39.
40.
41.
42.
43.
44.
45.
46.
47.
48.
49.
50.
51.
52.
53.
54.
55.
56.
OBJ: 9.2.2 Apply the rules for naming and writing formulas for compounds with polyatomic ions. | 9.5.2
Apply the rules for naming chemical compounds by using a flowchart.
STA: 5.6.A.5
ANS: B
PTS: 1
DIF: L3
REF: p. 264 | p. 277
OBJ: 9.5.3 Apply the rules for writing chemical formulas by using a flowchart.
STA: 5.6.A.5
ANS: B
PTS: 1
DIF: L3
REF: p. 290 | p. 291
OBJ: 10.1.2 Relate Avogadro’s number to a mole of a substance.
ANS: A
PTS: 1
DIF: L3
REF: p. 301
OBJ: 10.2.2 Identify the volume of a quantity of gas at STP.
ANS: D
PTS: 1
DIF: L3
REF: p. 301
OBJ: 10.2.2 Identify the volume of a quantity of gas at STP.
ANS: C
PTS: 1
DIF: L3
REF: p. 302
OBJ: 10.2.2 Identify the volume of a quantity of gas at STP.
ANS: D
PTS: 1
DIF: L3
REF: p. 307
OBJ: 10.3.1 Describe how to calculate the percent by mass of an element in a compound.
ANS: D
PTS: 1
DIF: L3
REF: p. 307
OBJ: 10.3.1 Describe how to calculate the percent by mass of an element in a compound.
ANS: D
PTS: 1
DIF: L3
REF: p. 310
OBJ: 10.3.2 Interpret an empirical formula.
ANS: C
PTS: 1
DIF: L2
REF: p. 101
OBJ: 4.1.1 Describe Democritus's ideas about atoms.
STA: 5.2.B.1
ANS: C
PTS: 1
DIF: L2
REF: p. 101
OBJ: 4.1.1 Describe Democritus's ideas about atoms.
STA: 5.2.B.1
ANS: C
PTS: 1
DIF: L2
REF: p. 102
OBJ: 4.1.2 Explain Dalton's atomic theory.
STA: 5.2.B.1
ANS: B
PTS: 1
DIF: L2
REF: p. 102
OBJ: 4.1.2 Explain Dalton's atomic theory.
STA: 5.2.B.1
ANS: A
PTS: 1
DIF: L2
REF: p. 104
OBJ: 4.1.2 Explain Dalton's atomic theory.
STA: 5.2.B.1
ANS: B
PTS: 1
DIF: L2
REF: p. 103
OBJ: 4.1.3 Identify the special instruments necessary to observe individual atoms.
STA: 5.6.A.1
ANS: C
PTS: 1
DIF: L2
REF: p. 103
OBJ: 4.1.3 Identify the special instruments necessary to observe individual atoms.
STA: 5.6.A.1
ANS: D
PTS: 1
DIF: L2
REF: p. 105
OBJ: 4.2.1 Identify three types of subatomic particle.
STA: 5.2.B.1
ANS: D
PTS: 1
DIF: L2
REF: p. 104 | p. 105 | p. 106
OBJ: 4.2.1 Identify three types of subatomic particle.
STA: 5.6.A.1
ANS: B
PTS: 1
DIF: L2
REF: p. 105
OBJ: 4.2.1 Identify three types of subatomic particle.
STA: 5.2.B.1
ANS: A
PTS: 1
DIF: L2
REF: p. 105
OBJ: 4.2.1 Identify three types of subatomic particle.
STA: 5.6.A.1
ANS: C
PTS: 1
DIF: L2
REF: p. 106 | p. 107
OBJ: 4.2.1 Identify three types of subatomic particle. | 4.2.2 Describe the structure of atoms according to
the Rutherford model.
STA:
5.6.A.1
ANS: C
PTS: 1
DIF: L2
REF: p. 108
OBJ: 4.2.2 Describe the structure of atoms according to the Rutherford model.
STA: 5.2.B.1
57. ANS: A
PTS: 1
DIF: L2
REF: p. 107 | p. 108
OBJ: 4.2.2 Describe the structure of atoms according to the Rutherford model.
STA: 5.6.A.1
58. ANS: A
PTS: 1
DIF: L2
REF: p. 110
OBJ: 4.2.1 Identify three types of subatomic particle. | 4.3.1 Explain what makes elements and isotopes
different from each other.
STA: 5.6.A.2
59. ANS: B
PTS: 1
DIF: L2
REF: p. 111
OBJ: 4.3.1 Explain what makes elements and isotopes different from each other.
STA: 5.6.A.1
60. ANS: B
PTS: 1
DIF: L2
REF: p. 111
OBJ: 4.3.1 Explain what makes elements and isotopes different from each other. | 4.3.4 Explain why
chemists use the periodic table.
STA: 5.6.A.5
61. ANS: A
PTS: 1
DIF: L2
REF: p. 111
OBJ: 4.3.1 Explain what makes elements and isotopes different from each other.
STA: 5.6.A.1
62. ANS: B
PTS: 1
DIF: L2
REF: p. 110 | p. 112 | p. 113
OBJ: 4.3.1 Explain what makes elements and isotopes different from each other.
63. ANS: C
PTS: 1
DIF: L2
REF: p. 112
OBJ: 4.3.1 Explain what makes elements and isotopes different from each other.
64. ANS: D
PTS: 1
DIF: L2
REF: p. 113
OBJ: 4.3.1 Explain what makes elements and isotopes different from each other.
65. ANS: C
PTS: 1
DIF: L2
REF: p. 111
OBJ: 4.3.2 Calculate the number of neutrons in an atom.
STA: 5.6.A.1
66. ANS: D
PTS: 1
DIF: L2
REF: p. 112 | p. 113
OBJ: 4.3.2 Calculate the number of neutrons in an atom.
STA: 5.6.A.1
67. ANS: B
PTS: 1
DIF: L2
REF: p. 114
OBJ: 4.3.3 Calculate the atomic mass of an element.
68. ANS: B
PTS: 1
DIF: L2
REF: p. 114
OBJ: 4.3.3 Calculate the atomic mass of an element.
69. ANS: B
PTS: 1
DIF: L2
REF: p. 115
OBJ: 4.3.3 Calculate the atomic mass of an element.
70. ANS: D
PTS: 1
DIF: L2
REF: p. 115
OBJ: 4.3.3 Calculate the atomic mass of an element.
71. ANS: D
PTS: 1
DIF: L2
REF: p. 116
OBJ: 4.3.3 Calculate the atomic mass of an element.
72. ANS: C
PTS: 1
DIF: L2
REF: p. 131
OBJ: 5.1.3 Describe the energies and positions of electrons according to the quantum mechanical model.
STA: 5.6.A.8
73. ANS: A
PTS: 1
DIF: L2
REF: p. 157
OBJ: 6.1.1 Explain how elements are organized in a periodic table.
STA: 5.6.A.5
74. ANS: C
PTS: 1
DIF: L2
REF: p. 157
OBJ: 6.1.1 Explain how elements are organized in a periodic table.
STA: 5.6.A.5
75. ANS: D
PTS: 1
DIF: L2
REF: p. 158
OBJ: 6.1.3 Identify three broad classes of elements.
STA: 5.6.A.5
76. ANS: A
PTS: 1
DIF: L2
REF: p. 157
OBJ: 6.2.1 Describe the information in a periodic table.
STA: 5.6.A.1
77. ANS: C
PTS: 1
DIF: L2
REF: p. 164
OBJ: 6.2.2 Classify elements based on electron configuration. STA: 5.6.A.5
78. ANS: A
PTS: 1
DIF: L2
REF: p. 164
OBJ: 6.2.2 Classify elements based on electron configuration. STA: 5.6.A.5
79. ANS: A
PTS: 1
DIF: L2
REF: p. 166
OBJ: 6.2.2 Classify elements based on electron configuration. STA: 5.6.A.5
80. ANS: D
PTS: 1
DIF: L2
REF: p. 164
OBJ: 6.2.2 Classify elements based on electron configuration. STA: 5.6.A.5 | 5.6.A.3
81. ANS: B
PTS: 1
DIF: L2
REF: p. 164
OBJ: 6.2.2 Classify elements based on electron configuration. STA: 5.6.A.5
82. ANS: B
PTS: 1
DIF: L2
REF: p. 164
OBJ: 6.2.2 Classify elements based on electron configuration. | 6.2.3 Distinguish representative elements
and transition metals.
STA: 5.6.A.5
83. ANS: A
PTS: 1
DIF: L2
REF: p. 164
OBJ: 6.2.3 Distinguish representative elements and transition metals.
STA: 5.6.A.5
84. ANS: A
PTS: 1
DIF: L2
REF: p. 171
OBJ: 6.3.1 Describe trends among elements for atomic size.
STA: 5.6.A.5
85. ANS: A
PTS: 1
DIF: L2
REF: p. 171
OBJ: 6.3.1 Describe trends among elements for atomic size.
STA: 5.6.A.5
86. ANS: D
PTS: 1
DIF: L2
REF: p. 171
OBJ: 6.3.1 Describe trends among elements for atomic size.
STA: 5.6.A.5
87. ANS: B
PTS: 1
DIF: L2
REF: p. 171
OBJ: 6.3.1 Describe trends among elements for atomic size.
STA: 5.6.A.5
88. ANS: B
PTS: 1
DIF: L2
REF: p. 172
OBJ: 6.3.2 Explain how ions form.
STA: 5.6.A.4 | 5.6.A.5
89. ANS: D
PTS: 1
DIF: L2
REF: p. 162 | p. 163 | p. 172 | p. 176
OBJ: 6.3.2 Explain how ions form.
STA: 5.6.A.5
90. ANS: D
PTS: 1
DIF: L2
REF: p. 172
OBJ: 6.3.2 Explain how ions form.
STA: 5.6.A.4
91. ANS: A
PTS: 1
DIF: L2
REF: p. 173
OBJ: 6.3.2 Explain how ions form.
STA: 5.6.A.4
92. ANS: A
PTS: 1
DIF: L2
REF: p. 175
OBJ: 6.3.3 Describe periodic trends for first ionization energy, ionic size, and electronegativity.
STA: 5.6.A.4 | 5.6.A.5
93. ANS: B
PTS: 1
DIF: L2
REF: p. 173
OBJ: 6.3.3 Describe periodic trends for first ionization energy, ionic size, and electronegativity.
STA: 5.6.A.5 | 5.6.A.8
94. ANS: C
PTS: 1
DIF: L2
REF: p. 173
OBJ: 6.3.3 Describe periodic trends for first ionization energy, ionic size, and electronegativity.
STA: 5.6.A.5 | 5.6.A.8
95. ANS: A
PTS: 1
DIF: L2
REF: p. 174
OBJ: 6.3.1 Describe trends among elements for atomic size. | 6.3.3 Describe periodic trends for first
ionization energy, ionic size, and electronegativity.
STA: 5.6.A.5 | 5.6.A.8
96. ANS: C
PTS: 1
DIF: L2
REF: p. 173
OBJ: 6.3.3 Describe periodic trends for first ionization energy, ionic size, and electronegativity.
STA: 5.6.A.5 | 5.6.A.8
97. ANS: A
PTS: 1
DIF: L2
REF: p. 177
OBJ: 6.3.3 Describe periodic trends for first ionization energy, ionic size, and electronegativity.
STA: 5.6.A.5
98. ANS: D
PTS: 1
DIF: L2
REF: p. 177
OBJ: 6.3.3 Describe periodic trends for first ionization energy, ionic size, and electronegativity.
STA: 5.6.A.5
99. ANS: B
PTS: 1
DIF: L2
REF: p. 177 | p. 178
OBJ: 6.3.3 Describe periodic trends for first ionization energy, ionic size, and electronegativity.
STA: 5.6.A.5
100. ANS: C
PTS: 1
DIF: L2
REF: p. 178
OBJ: 6.3.1 Describe trends among elements for atomic size. | 6.3.3 Describe periodic trends for first
ionization energy, ionic size, and electronegativity.
STA: 5.6.A.5
101. ANS: A
PTS: 1
DIF: L2
REF: p. 172 | p. 176
OBJ: 6.3.3 Describe periodic trends for first ionization energy, ionic size, and electronegativity.
STA: 5.6.A.5
102. ANS: A
PTS: 1
DIF: L2
REF: p. 188 | p. 189
OBJ: 7.1.1 Determine the number of valence electrons in an atom of a representative element.
STA: 5.6.A.4
103. ANS: D
PTS: 1
DIF: L2
REF: p. 190
OBJ: 7.1.1 Determine the number of valence electrons in an atom of a representative element.
STA: 5.6.A.4
104. ANS: A
PTS: 1
DIF: L2
REF: p. 188
OBJ: 7.1.2 Explain how the octet rule applies to atoms of metallic and nonmetallic elements.
STA: 5.6.A.4
105. ANS: A
PTS: 1
DIF: L2
REF: p. 190
OBJ: 7.1.3 Describe how cations form. STA: 5.6.A.4
106. ANS: B
PTS: 1
DIF: L2
REF: p. 190
OBJ: 7.1.3 Describe how cations form. STA: 5.6.A.4
107. ANS: C
PTS: 1
DIF: L2
REF: p. 190
OBJ: 7.1.3 Describe how cations form. STA: 5.6.A.4
108. ANS: B
PTS: 1
DIF: L2
REF: p. 192
OBJ: 7.1.4 Explain how anions form.
STA: 5.6.A.4
109. ANS: A
PTS: 1
DIF: L2
REF: p. 192
OBJ: 7.1.4 Explain how anions form.
STA: 5.6.A.4
110. ANS: A
PTS: 1
DIF: L2
REF: p. 194
OBJ: 7.2.1 Explain the electrical charge of an ionic compound.
STA: 5.6.A.3 | 5.6.A.4
111. ANS: C
PTS: 1
DIF: L2
REF: p. 194
OBJ: 7.2.1 Explain the electrical charge of an ionic compound.
STA: 5.6.A.3 | 5.6.A.4
112. ANS: C
PTS: 1
DIF: L2
REF: p. 195
OBJ: 7.2.1 Explain the electrical charge of an ionic compound.
113. ANS: D
PTS: 1
DIF: L2
REF: p. 195
OBJ: 7.2.1 Explain the electrical charge of an ionic compound.
114. ANS: B
PTS: 1
DIF: L2
REF: p. 192 | p. 195
OBJ: 7.2.1 Explain the electrical charge of an ionic compound.
STA: 5.6.A.3
115. ANS: B
PTS: 1
DIF: L2
REF: p. 192 | p. 195
OBJ: 7.2.1 Explain the electrical charge of an ionic compound.
STA: 5.6.A.3
116. ANS: D
PTS: 1
DIF: L2
REF: p. 203
OBJ: 7.3.3 Explain the importance of alloys.
STA: 5.6.A.3
117. ANS: A
PTS: 1
DIF: L2
REF: p. 192 | p. 194
OBJ: 7.2.1 Explain the electrical charge of an ionic compound.
STA: 5.6.A.3
118. ANS: D
PTS: 1
DIF: L2
REF: p. 244
OBJ: 8.1.1 Distinguish between the melting points and boiling points of molecular compounds and ionic
compounds.
STA: 5.6.A.7
119. ANS: C
PTS: 1
DIF: L2
REF: p. 217
OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet
rule. STA:
5.6.A.4
120. ANS: B
PTS: 1
DIF: L2
REF: p. 217
OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet
rule. STA:
5.6.A.4
121. ANS: D
PTS: 1
DIF: L2
REF: p. 221
OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet
rule. STA:
5.6.A.4
122. ANS: B
PTS: 1
DIF: L2
REF: p. 218
OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet
rule.
123. ANS: A
PTS: 1
DIF: L2
REF: p. 221
OBJ: 8.2.3 Describe how atoms form double or triple covalent bonds.
STA: 5.6.A.4
124. ANS: B
PTS: 1
DIF: L2
REF: p. 222
OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet
rule. | 8.2.4 Distinguish between a covalent bond and a coordinate covalent bond and describe how the
strength of a covalent bond is related to its bond dissociation energy.
STA: 5.6.A.4
125. ANS: C
PTS: 1
DIF: L2
REF: p. 223
OBJ: 8.2.4 Distinguish between a covalent bond and a coordinate covalent bond and describe how the
strength of a covalent bond is related to its bond dissociation energy.
STA: 5.6.A.4
126. ANS: D
PTS: 1
DIF: L2
REF: p. 223
OBJ: 8.2.4 Distinguish between a covalent bond and a coordinate covalent bond and describe how the
strength of a covalent bond is related to its bond dissociation energy.
STA: 5.6.A.4
127. ANS: A
PTS: 1
DIF: L2
REF: p. 225
OBJ: 8.2.4 Distinguish between a covalent bond and a coordinate covalent bond and describe how the
strength of a covalent bond is related to its bond dissociation energy.
STA: 5.6.A.4
128. ANS: B
PTS: 1
DIF: L2
REF: p. 226
OBJ: 8.2.5 Describe how oxygen atoms are bonded in ozone. STA: 5.6.A.4
129. ANS: D
PTS: 1
DIF: L2
REF: p. 229
OBJ: 8.2.5 Describe how oxygen atoms are bonded in ozone.
130. ANS: D
PTS: 1
DIF: L2
REF: p. 230
OBJ: 8.3.1 Describe the relationship between atomic and molecular orbitals.
131. ANS: C
PTS: 1
DIF: L2
REF: p. 231
OBJ: 8.3.1 Describe the relationship between atomic and molecular orbitals.
STA: 5.6.A.4
132. ANS: B
PTS: 1
DIF: L2
REF: p. 231
OBJ: 8.3.1 Describe the relationship between atomic and molecular orbitals.
133. ANS: D
PTS: 1
DIF: L2
REF: p. 230 | p. 231
OBJ: 8.3.1 Describe the relationship between atomic and molecular orbitals.
STA: 5.6.A.4
134. ANS: D
PTS: 1
DIF: L2
REF: p. 231
OBJ: 8.3.1 Describe the relationship between atomic and molecular orbitals.
STA: 5.6.A.4
135. ANS: A
PTS: 1
DIF: L2
REF: p. 233
OBJ: 8.3.2 Describe how VSEPR theory helps predict the shapes of molecules.
136. ANS: C
PTS: 1
DIF: L2
REF: p. 234
OBJ: 8.3.3 Identify the ways in which orbital hybridization is useful in describing molecules.
137. ANS: D
PTS: 1
DIF: L2
REF: p. 235
OBJ: 8.3.3 Identify the ways in which orbital hybridization is useful in describing molecules.
STA: 5.6.A.4
138. ANS: C
PTS: 1
DIF: L2
REF: p. 235
OBJ: 8.3.3 Identify the ways in which orbital hybridization is useful in describing molecules.
STA: 5.6.A.4
139. ANS: B
PTS: 1
DIF: L2
REF: p. 235
OBJ: 8.3.3 Identify the ways in which orbital hybridization is useful in describing molecules.
STA: 5.6.A.4
140. ANS: C
PTS: 1
DIF: L2
REF: p. 238 | p. 239
OBJ: 8.4.1 Describe how electronegativity values determine the charge distribution in a polar molecule.
STA: 5.6.A.4
141. ANS: C
PTS: 1
DIF: L2
REF: p. 238 | p. 239
OBJ: 8.1.1 Distinguish between the melting points and boiling points of molecular compounds and ionic
compounds. | 8.4.1 Describe how electronegativity values determine the charge distribution in a polar
molecule.
STA: 5.6.A.4
142. ANS: D
PTS: 1
DIF: L2
REF: p. 241
OBJ: 8.4.3 Evaluate the strength of intermolecular attractions compared with the strength of ionic and
covalent bonds.
STA: 5.6.A.4
143. ANS: A
PTS: 1
DIF: L2
REF: p. 241
OBJ: 8.4.1 Describe how electronegativity values determine the charge distribution in a polar molecule. |
8.4.3 Evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent
bonds. STA:
5.6.A.4
144. ANS: A
PTS: 1
DIF: L2
REF: p. 254 | p. 255
OBJ: 9.1.1 Identify the charges of monatomic ions by using the periodic table, and name the ions.
STA: 5.6.A.5
145. ANS: C
PTS: 1
DIF: L2
REF: p. 254 | p. 255
OBJ: 9.1.1 Identify the charges of monatomic ions by using the periodic table, and name the ions.
STA: 5.6.A.5
146. ANS: B
PTS: 1
DIF: L2
REF: p. 254
OBJ: 9.1.1 Identify the charges of monatomic ions by using the periodic table, and name the ions.
STA: 5.6.A.5
147. ANS: D
PTS: 1
DIF: L2
REF: p. 253 | p. 254
OBJ: 9.1.1 Identify the charges of monatomic ions by using the periodic table, and name the ions.
STA: 5.6.A.5
148. ANS: C
PTS: 1
DIF: L2
REF: p. 255
OBJ: 9.1.1 Identify the charges of monatomic ions by using the periodic table, and name the ions.
STA: 5.6.A.5
149. ANS: C
PTS: 1
DIF: L2
REF: p. 254 | p. 257
OBJ: 9.1.1 Identify the charges of monatomic ions by using the periodic table, and name the ions. | 9.1.2
Define a polyatomic ion and write the names and formulas of the most common polyatomic ions.
STA: 5.6.A.5
150. ANS: D
PTS: 1
DIF: L2
REF: p. 257
OBJ: 9.1.2 Define a polyatomic ion and write the names and formulas of the most common polyatomic
ions. STA:
5.6.A.5
151. ANS: D
PTS: 1
DIF: L2
REF: p. 257
OBJ: 9.1.2 Define a polyatomic ion and write the names and formulas of the most common polyatomic
ions. STA:
5.6.A.5
152. ANS: A
PTS: 1
DIF: L2
REF: p. 261
OBJ: 9.2.1 Apply the rules for naming and writing formulas for binary ionic compounds.
STA: 5.6.A.1
153. ANS: B
PTS: 1
DIF: L2
REF: p. 262
OBJ: 9.2.1 Apply the rules for naming and writing formulas for binary ionic compounds.
STA: 5.6.A.5
154. ANS: A
PTS: 1
DIF: L2
REF: p. 253 | p. 254 | p. 262
OBJ: 9.2.1 Apply the rules for naming and writing formulas for binary ionic compounds.
STA: 5.6.A.2
155. ANS: D
PTS: 1
DIF: L2
REF: p. 262
OBJ: 9.2.1 Apply the rules for naming and writing formulas for binary ionic compounds.
STA: 5.6.A.5
156. ANS: B
PTS: 1
DIF: L2
REF: p. 262
OBJ: 9.2.1 Apply the rules for naming and writing formulas for binary ionic compounds.
STA: 5.6.A.5
157. ANS: A
PTS: 1
DIF: L2
REF: p. 262
OBJ: 9.2.1 Apply the rules for naming and writing formulas for binary ionic compounds.
STA: 5.6.A.5
158. ANS: D
PTS: 1
DIF: L2
REF: p. 262 | p. 263
OBJ: 9.2.1 Apply the rules for naming and writing formulas for binary ionic compounds.
STA: 5.6.A.5
159. ANS: A
PTS: 1
DIF: L2
REF: p. 262 | p. 263
OBJ: 9.2.1 Apply the rules for naming and writing formulas for binary ionic compounds.
STA: 5.6.A.5
160. ANS: C
PTS: 1
DIF: L2
REF: p. 261 | p. 262
OBJ: 9.2.1 Apply the rules for naming and writing formulas for binary ionic compounds.
STA: 5.6.A.5
161. ANS: C
PTS: 1
DIF: L2
REF: p. 257 | p. 261 | p. 262
OBJ: 9.2.2 Apply the rules for naming and writing formulas for compounds with polyatomic ions.
STA: 5.6.A.5
162. ANS: B
PTS: 1
DIF: L2
REF: p. 264 | p. 265 | p. 266
OBJ: 9.1.3 Identify the two common endings for the names of most polyatomic ions. | 9.2.2 Apply the rules
for naming and writing formulas for compounds with polyatomic ions.
STA: 5.6.A.5
163. ANS: C
PTS: 1
DIF: L2
REF: p. 264 | p. 277
OBJ: 9.2.2 Apply the rules for naming and writing formulas for compounds with polyatomic ions.
STA: 5.6.A.5
164. ANS: C
PTS: 1
DIF: L2
REF: p. 268
OBJ: 9.3.1 Interpret the prefixes in the names of molecular compounds in terms of their chemical formulas.
STA: 5.6.A.5
165. ANS: B
PTS: 1
DIF: L2
REF: p. 269
OBJ: 9.3.2 Apply the rules for naming and writing formulas for binary molecular compounds.
STA: 5.6.A.5
166. ANS: B
PTS: 1
DIF: L2
REF: p. 268
OBJ: 9.3.2 Apply the rules for naming and writing formulas for binary molecular compounds.
STA: 5.6.A.5
167. ANS: C
PTS: 1
DIF: L2
REF: p. 269
OBJ: 9.3.2 Apply the rules for naming and writing formulas for binary molecular compounds.
STA: 5.6.A.5
168. ANS: A
PTS: 1
DIF: L2
REF: p. 272
OBJ: 9.4.1 Apply three rules for naming acids.
STA: 5.6.A.5
169. ANS: C
PTS: 1
DIF: L2
REF: p. 272
OBJ: 9.4.1 Apply three rules for naming acids.
STA: 5.6.A.5
170. ANS: D
PTS: 1
DIF: L2
REF: p. 272
OBJ: 9.4.1 Apply three rules for naming acids.
STA: 5.6.A.5
171. ANS: A
PTS: 1
DIF: L2
REF: p. 272
OBJ: 9.4.1 Apply three rules for naming acids.
STA: 5.6.A.5
172. ANS: B
PTS: 1
DIF: L2
REF: p. 272
OBJ: 9.4.2 Apply the rules in reverse to write formulas of acids.
STA: 5.6.A.5
173. ANS: B
PTS: 1
DIF: L2
REF: p. 272
OBJ: 9.4.2 Apply the rules in reverse to write formulas of acids.
STA: 5.6.A.5
174. ANS: D
PTS: 1
DIF: L2
REF: p. 274
OBJ: 9.5.1 Define the laws of definition proportions and multiple proportions.
175. ANS: D
PTS: 1
DIF: L2
REF: p. 270 | p. 278
OBJ: 9.3.2 Apply the rules for naming and writing formulas for binary molecular compounds. | 9.5.2 Apply
the rules for naming chemical compounds by using a flowchart.
STA: 5.6.A.5
176. ANS: D
PTS: 1
DIF: L2
REF: p. 261 | p. 262 | p. 277
OBJ: 9.2.1 Apply the rules for naming and writing formulas for binary ionic compounds. | 9.5.2 Apply the
rules for naming chemical compounds by using a flowchart.
STA: 5.6.A.5
177. ANS: B
PTS: 1
DIF: L2
REF: p. 271 | p. 276 | p. 277
OBJ: 9.4.1 Apply three rules for naming acids. | 9.5.2 Apply the rules for naming chemical compounds by
using a flowchart.
178. ANS: B
PTS: 1
DIF: L2
REF: p. 269 | p. 277
OBJ: 9.3.2 Apply the rules for naming and writing formulas for binary molecular compounds. | 9.5.3 Apply
the rules for writing chemical formulas by using a flowchart.
STA: 5.6.A.5
179. ANS: B
PTS: 1
DIF: L2
REF: p. 290 | p. 291
OBJ: 10.1.2 Relate Avogadro’s number to a mole of a substance.
180. ANS: A
PTS: 1
DIF: L2
REF: p. 290 | p. 291
OBJ: 10.1.2 Relate Avogadro’s number to a mole of a substance.
181. ANS: D
PTS: 1
DIF: L2
REF: p. 291 | p. 292
OBJ: 10.1.2 Relate Avogadro’s number to a mole of a substance.
182. ANS: D
PTS: 1
DIF: L2
REF: p. 291 | p. 292
OBJ: 10.1.2 Relate Avogadro’s number to a mole of a substance.
183. ANS: C
PTS: 1
DIF: L2
REF: p. 291 | p. 292
OBJ: 10.1.2 Relate Avogadro’s number to a mole of a substance.
184. ANS: A
PTS: 1
DIF: L2
REF: p. 294
OBJ: 10.1.3 Distiguish between the atomic mass of an element and its molar mass.
185. ANS: B
PTS: 1
DIF: L2
REF: p. 294
OBJ: 10.1.3 Distiguish between the atomic mass of an element and its molar mass.
186. ANS: D
PTS: 1
DIF: L2
REF: p. 295 | p. 296
OBJ: 10.1.4 Describe how the mass of a mole of a compound is calculated.
187. ANS: C
PTS: 1
DIF: L2
REF: p. 295 | p. 296
OBJ: 10.1.4 Describe how the mass of a mole of a compound is calculated.
188. ANS: D
PTS: 1
DIF: L2
REF: p. 295 | p. 296
OBJ: 10.1.4 Describe how the mass of a mole of a compound is calculated.
189. ANS: D
PTS: 1
DIF: L2
REF: p. 297 | p. 298
OBJ: 10.2.1 Describe how to convert the mass of a substance to the number of moles of a substance, and
moles to mass.
190. ANS: C
PTS: 1
DIF: L2
REF: p. 299
OBJ: 10.2.1 Describe how to convert the mass of a substance to the number of moles of a substance, and
moles to mass.
191. ANS: B
PTS: 1
DIF: L2
REF: p. 299
OBJ: 10.2.1 Describe how to convert the mass of a substance to the number of moles of a substance, and
moles to mass.
192. ANS: A
PTS: 1
DIF: L2
REF: p. 299
OBJ: 10.2.1 Describe how to convert the mass of a substance to the number of moles of a substance, and
moles to mass.
193. ANS: C
PTS: 1
DIF: L2
REF: p. 297
OBJ: 10.2.1 Describe how to convert the mass of a substance to the number of moles of a substance, and
moles to mass.
194. ANS: C
PTS: 1
DIF: L2
REF: p. 298
OBJ: 10.2.1 Describe how to convert the mass of a substance to the number of moles of a substance, and
moles to mass.
195. ANS: D
PTS: 1
DIF: L2
REF: p. 298
OBJ: 10.2.1 Describe how to convert the mass of a substance to the number of moles of a substance, and
moles to mass.
196. ANS: B
PTS: 1
DIF: L2
REF: p. 301
OBJ: 10.2.2 Identify the volume of a quantity of gas at STP.
197. ANS: C
PTS: 1
DIF: L2
REF: p. 301
OBJ: 10.2.2 Identify the volume of a quantity of gas at STP.
198. ANS: B
PTS: 1
DIF: L2
REF: p. 301
OBJ: 10.2.2 Identify the volume of a quantity of gas at STP.
199. ANS: B
PTS: 1
DIF: L2
REF: p. 302
OBJ: 10.2.2 Identify the volume of a quantity of gas at STP.
200. ANS: C
PTS: 1
DIF: L2
REF: p. 302
OBJ: 10.2.2 Identify the volume of a quantity of gas at STP.
201. ANS: A
PTS: 1
DIF: L2
REF: p. 300
OBJ: 10.2.2 Identify the volume of a quantity of gas at STP.
202. ANS: D
PTS: 1
DIF: L2
REF: p. 302
OBJ: 10.2.2 Identify the volume of a quantity of gas at STP.
203. ANS: C
PTS: 1
DIF: L2
REF: p. 305 | p. 306
OBJ: 10.3.1 Describe how to calculate the percent by mass of an element in a compound.
204. ANS: C
PTS: 1
DIF: L2
REF: p. 307
OBJ:
205. ANS:
OBJ:
206. ANS:
OBJ:
207. ANS:
OBJ:
208. ANS:
OBJ:
209. ANS:
OBJ:
210. ANS:
OBJ:
211. ANS:
OBJ:
212. ANS:
OBJ:
213. ANS:
OBJ:
STA:
214. ANS:
OBJ:
STA:
215. ANS:
OBJ:
STA:
216. ANS:
OBJ:
217. ANS:
OBJ:
218. ANS:
OBJ:
219. ANS:
OBJ:
220. ANS:
OBJ:
221. ANS:
OBJ:
222. ANS:
OBJ:
223. ANS:
OBJ:
224. ANS:
OBJ:
225. ANS:
OBJ:
226. ANS:
OBJ:
227. ANS:
10.3.1 Describe how to calculate the percent by mass of an element in a compound.
D
PTS: 1
DIF: L2
REF: p. 305 | p. 306
10.3.1 Describe how to calculate the percent by mass of an element in a compound.
D
PTS: 1
DIF: L2
REF: p. 307
10.3.1 Describe how to calculate the percent by mass of an element in a compound.
B
PTS: 1
DIF: L2
REF: p. 307
10.3.1 Describe how to calculate the percent by mass of an element in a compound.
B
PTS: 1
DIF: L2
REF: p. 307
10.3.1 Describe how to calculate the percent by mass of an element in a compound.
C
PTS: 1
DIF: L2
REF: p. 310
10.3.2 Interpret an empirical formula.
C
PTS: 1
DIF: L2
REF: p. 310
10.3.2 Interpret an empirical formula.
B
PTS: 1
DIF: L2
REF: p. 312
10.3.3 Distinguish between empirical and molecular formulas.
C
PTS: 1
DIF: L2
REF: p. 323
11.1.2 Describe how to write a skeleton equation
STA: 5.6.B.1
C
PTS: 1
DIF: L2
REF: p. 324 | p. 325
11.1.3 Describe the steps for writing a balanced chemical equation.
5.6.B.1
D
PTS: 1
DIF: L2
REF: p. 324 | p. 327
11.1.3 Describe the steps for writing a balanced chemical equation.
5.6.B.1
C
PTS: 1
DIF: L2
REF: p. 327
11.1.3 Describe the steps for writing a balanced chemical equation.
5.6.B.1
B
PTS: 1
DIF: L2
REF: p. 332
11.2.1 Describe the five general types of reactions.
STA: 5.6.B.1
C
PTS: 1
DIF: L2
REF: p. 334 | p. 335
11.2.1 Describe the five general types of reactions.
STA: 5.6.B.1
D
PTS: 1
DIF: L2
REF: p. 336 | p. 337
11.2.1 Describe the five general types of reactions.
STA: 5.6.B.1
A
PTS: 1
DIF: L2
REF: p. 332
11.2.1 Describe the five general types of reactions.
STA: 5.6.B.1
C
PTS: 1
DIF: L2
REF: p. 333 | p. 334
11.2.1 Describe the five general types of reactions.
STA: 5.6.B.1
A
PTS: 1
DIF: L2
REF: p. 333
11.2.1 Describe the five general types of reactions.
STA: 5.6.B.1
D
PTS: 1
DIF: L2
REF: p. 333 | p. 334
11.2.1 Describe the five general types of reactions.
STA: 5.6.B.1
D
PTS: 1
DIF: L2
REF: p. 334 | p. 335
11.2.1 Describe the five general types of reactions.
STA: 5.6.B.1
C
PTS: 1
DIF: L2
REF: p. 334 | p. 335
11.2.1 Describe the five general types of reactions.
STA: 5.6.B.1
A
PTS: 1
DIF: L2
REF: p. 334 | p. 335
11.2.1 Describe the five general types of reactions.
STA: 5.6.B.1
A
PTS: 1
DIF: L2
REF: p. 336 | p. 337
11.2.1 Describe the five general types of reactions.
STA: 5.6.B.1
B
PTS: 1
DIF: L2
REF: p. 336
OBJ:
228. ANS:
OBJ:
STA:
229. ANS:
OBJ:
STA:
230. ANS:
OBJ:
STA:
11.2.1 Describe the five general types of reactions.
STA: 5.6.B.1
D
PTS: 1
DIF: L2
REF: p. 330 | p. 337
11.2.2 Predict the products of the five general types of reactions.
5.6.B.1
B
PTS: 1
DIF: L2
REF: p. 333 | p. 334
11.2.2 Predict the products of the five general types of reactions.
5.6.B.1
B
PTS: 1
DIF: L2
REF: p. 334 | p. 344
11.3.2 Predict the formation of a precipitate in a double-replacement reaction.
5.6.B.1