Review 4-9
Matching
Match each item with the correct statement below.
a. atomic orbital
d. ground state
b. aufbau principle
e. Pauli exclusion principle
c. electron configuration
f. Heisenberg uncertainty principle
____
1. region of high probability of finding an electron
____
2. states the impossibility of knowing both velocity and position of a moving particle at the same time
____
3. lowest energy level
____
4. tendency of electrons to enter orbitals of lowest energy first
____
5. arrangement of electrons around atomic nucleus
____
6. each orbital has at most two electrons
Match each item with the correct statement below.
a. atomic emission spectrum
d. photon
b. frequency
e. quantum
c. wavelength
f. spectrum
____
7. discrete bundle of electromagnetic energy
____
8. energy needed to move an electron from one energy level to another
____
9. number of wave cycles passing a point per unit of time
____ 10. distance between wave crests
____ 11. separation of light into different wavelengths
____ 12. frequencies of light emitted by an element
Match each item with the correct statement below.
a. coordinate covalent bond
d. single covalent bond
b. double covalent bond
e. polar bond
c. structural formula
f. hydrogen bond
____ 13. a depiction of the arrangement of atoms in molecules and polyatomic ions
____ 14. a covalent bond in which only one pair of electrons is shared
____ 15. a covalent bond in which two pairs of electrons are shared
____ 16. a covalent bond in which the shared electron pair comes from only one of the atoms
____ 17. a covalent bond between two atoms of significantly different electronegativities
____ 18. a type of bond that is very important in determining the properties of water and of important biological
molecules such as proteins and DNA
Match each item with the correct statement below.
a. network solid
e. tetrahedral angle
b. bonding orbital
f. VSEPR theory
c. dipole interaction
g. sigma bond
d. bond dissociation energy
____ 19. energy needed to break a single bond between two covalently bonded atoms
____ 20. symmetrical bond along the axis between the two nuclei
____ 21. molecular orbital that can be occupied by two electrons of a covalent bond
____ 22. 109.5
____ 23. shapes adjust so valence-electron pairs are as far apart as possible
____ 24. attraction between polar molecules
____ 25. crystal in which all the atoms are covalently bonded to each other
Multiple Choice
Identify the choice that best completes the statement or answers the question.
____ 26. What is the maximum number of electrons in the second principal energy level?
a. 2
c. 18
b. 8
d. 32
____ 27. 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
____ 28. 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.
____ 29. How many unpaired electrons are in a sulfur atom (atomic number 16)?
a. 0
c. 2
b. 1
d. 3
____ 30. How many half-filled orbitals are in a bromine atom?
a. 1
c. 3
b. 2
d. 4
____ 31. 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
____ 32. What is the wavelength of an electromagnetic wave that travels at 3
MHz? (1 MHz = 1,000,000 Hz)
10 m/s and has a frequency of 60
a.
b. 60 MHz  300,000,000 m/s
c.
d. No answer can be determined from the information given.
____ 33. Which variable is directly proportional to frequency?
a. wavelength
c. position
b. velocity
d. energy
____ 34. 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.
____ 35. 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
____ 36. 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
____ 37. Of the elements Fe, Hg, U, and Te, which is a representative element?
a. Fe
c. U
b. Hg
d. Te
____ 38. 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
____ 39. Which of the following elements has the smallest atomic radius?
a. sulfur
c. selenium
b. chlorine
d. bromine
____ 40. 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
____ 41. 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
____ 42. Which of the following factors contributes to the increase in ionization energy from left to right across a
period?
a.
b.
c.
d.
an increase in the shielding effect
an increase in the size of the nucleus
an increase in the number of protons
fewer electrons in the highest occupied energy level
____ 43. 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
____ 44. Of the following elements, which one has the smallest first ionization energy?
a. boron
c. aluminum
b. carbon
d. silicon
____ 45. 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
____ 46. 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
____ 47. 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
____ 48. Which of the following covalent bonds is the most polar?
a. H—F
c. H—H
b. H—C
d. H—N
____ 49. 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
____ 50. 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
____ 51. 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.
____ 52. What is the formula for hydrosulfuric acid?
a. H S
b. H SO
c. HSO
d. H S
____ 53. What is the correct formula for barium chlorate?
a. Ba(ClO)
c. Ba(ClO )
b. Ba(ClO )
d. BaCl
____ 54. What is the correct formula for calcium dihydrogen phosphate?
a. CaH PO
c. Ca(H PO )
b. Ca H PO
d. Ca(H HPO )
____ 55. What is the correct name for Sn (PO ) ?
a. tritin diphosphate
b. tin(II) phosphate
c. tin(III) phosphate
d. tin(IV) phosphate
____ 56. In Bohr's model of the atom, where are the electrons and protons located?
a. The electrons move around the protons, which are at the center of the atom.
b. The electrons and protons move throughout the atom.
c. The electrons occupy fixed positions around the protons, which are at the center of the
atom.
d. The electrons and protons are located throughout the atom, but they are not free to move.
____ 57. In the Bohr model of the atom, an electron in an orbit has a fixed ____.
a. position
c. energy
b. color
d. size
____ 58. How does the energy of an electron change when the electron moves closer to the nucleus?
a. It decreases.
c. It stays the same.
b. It increases.
d. It doubles.
____ 59. The principal quantum number indicates what property of an electron?
a. position
c. energy level
b. speed
d. electron cloud shape
____ 60. What is the shape of the 3p atomic orbital?
a. sphere
b. dumbbell
c. bar
d. two perpendicular dumbbells
____ 61. How many energy sublevels are in the second principal energy level?
a. 1
c. 3
b. 2
d. 4
____ 62. What is the maximum number of f orbitals in any single energy level in an atom?
a. 1
c. 5
b. 3
d. 7
____ 63. What is the maximum number of d orbitals in a principal energy level?
a. 1
c. 3
b. 2
d. 5
____ 64. What is the maximum number of orbitals in the p sublevel?
a. 2
c. 4
b. 3
d. 5
____ 65. The shape (not the size) of an electron cloud is determined by the electron's ____.
a. energy sublevel
b. position
c. speed
d. principal quantum number
____ 66. The letter "p" in the symbol 4p indicates the ____.
a. spin of an electron
c. principle energy level
b. orbital shape
d. speed of an electron
____ 67. What types of atomic orbitals are in the third principal energy level?
a. s and p only
c. s, p, and d only
b. p and d only
d. s, p, d, and f
____ 68. What is the next atomic orbital in the series 1s, 2s, 2p, 3s, 3p?
a. 2d
c. 3f
b. 3d
d. 4s
____ 69. According to the aufbau principle, ____.
a. an orbital may be occupied by only two electrons
b. electrons in the same orbital must have opposite spins
c. electrons enter orbitals of highest energy first
d. electrons enter orbitals of lowest energy first
____ 70. What is the number of electrons in the outermost energy level of an oxygen atom?
a. 2
c. 6
b. 4
d. 8
____ 71. What is the electron configuration of potassium?
a. 1s 2s 2p 3s 3p 4s
c. 1s 2s 3s 3p 3d
b. 1s 2s 2p 3s 3p
d. 1s 2s 2p 3s 3p 4s
____ 72. What is the basis for exceptions to the aufbau diagram?
a. Filled and half-filled energy sublevels are more stable than partially-filled energy
sublevels.
b. Electron configurations are only probable.
c. Electron spins are more important than energy levels in determining electron
configuration.
d. Some elements have unusual atomic orbitals.
____ 73. Which electron configuration of the 4f energy sublevel is the most stable?
a. 4f
c. 4f
b. 4f
d. 4f
____ 74. How does the speed of visible light compare with the speed of gamma rays, when both speeds are measured
in a vacuum?
a. The speed of visible light is greater.
b. The speed of gamma rays is greater.
c. The speeds are the same.
d. No answer can be determined from the information given.
____ 75. Which color of visible light has the shortest wavelength?
a. yellow
c. blue
b. green
d. violet
____ 76. Which of the following electromagnetic waves have the highest frequencies?
a. ultraviolet light waves
b. X-rays
c. microwaves
d. gamma rays
____ 77. Which type of electromagnetic radiation includes the wavelength 10
a. gamma ray
c. radio wave
b. microwave
d. visible light
m?
____ 78. How are the frequency and wavelength of light related?
a. They are inversely proportional to each other.
b. Frequency equals wavelength divided by the speed of light.
c. Wavelength is determined by dividing frequency by the speed of light.
d. They are directly proportional to each other.
____ 79. The light given off by an electric discharge through sodium vapor is ____.
a. a continuous spectrum
c. of a single wavelength
b. an emission spectrum
d. white light
____ 80. Emission of light from an atom occurs when an electron ____.
a. drops from a higher to a lower energy level
b. jumps from a lower to a higher energy level
c. moves within its atomic orbital
d. falls into the nucleus
____ 81. As changes in energy levels of electrons increase, the frequencies of atomic line spectra they emit ____.
a. increase
c. remain the same
b. decrease
d. cannot be determined
____ 82. The atomic emission spectra of a sodium atom on Earth and of a sodium atom in the sun would be ____.
a. the same
b. different from each other
c. the same as those of several other elements
d. the same as each other only in the ultraviolet range
____ 83. What is the approximate energy of a photon having a frequency of 4
a. 3 10
c. 2 10
J
J
b. 3 10
d.
J
3 10 J
____ 84. What is the approximate frequency of a photon having an energy 5
a. 8 10 Hz
c. 3 10
Hz
b. 3 10
d.
Hz
1 10
Hz
10 Hz? (h = 6.6
10
J? (h = 6.6
10
10
J s)
J s)
____ 85. Which of the following quantum leaps would be associated with the greatest energy of emitted light?
a. n = 5 to n = 1
c. n = 2 to n = 5
b. n = 4 to n = 5
d. n = 5 to n = 4
____ 86. Bohr's model could only explain the spectra of which type of atoms?
a. single atoms with one electron
b. bonded atoms with one electron
c. single atoms with more than one electron
d. bonded atoms with more than one electron
____ 87. The quantum mechanical model of the atom ____.
a.
b.
c.
d.
defines the exact path of an electron around the nucleus
was proposed by Niels Bohr
involves the probability of finding an electron in a certain position
has many analogies in the visible world
____ 88. Who predicted that all matter can behave as waves as well as particles?
a. Albert Einstein
c. Max Planck
b. Erwin Schrodinger
d. Louis de Broglie
____ 89. According to the Heisenberg uncertainty principle, if the position of a moving particle is known, what other
quantity CANNOT be known?
a. mass
c. spin
b. charge
d. velocity
____ 90. How can the position of a particle be determined?
a. by analyzing its interactions with another particle
b. by measuring its velocity
c. by measuring its mass
d. by determining its charge
____ 91. The wavelike properties of electrons are useful in ____.
a. defining photons
b. writing electron configurations
c. magnifying objects
d. determining the velocity and position of a particle
____ 92. Each period in the periodic table corresponds to ____.
a. a principal energy level
c. an orbital
b. an energy sublevel
d. a suborbital
____ 93. The modern periodic table is arranged in order of increasing atomic ____.
a. mass
c. number
b. charge
d. radius
____ 94. Of the elements Pt, V, Li, and Kr, which is a nonmetal?
a. Pt
c. Li
b. V
d. Kr
____ 95. 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
____ 96. What element has the electron configuration 1s 2s 2p 3s 3p ?
a. nitrogen
c. silicon
b. selenium
d. silver
____ 97. 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.
____ 98. Elements that are characterized by the filling of p orbitals are classified as ____.
a. groups 3A through 8A
b. transition metals
c. inner transition metals
d. groups 1A and 2A
____ 99. 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
____ 100. 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
____ 101. 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.
____ 102. What are the Group 1A and Group 7A elements examples of?
a. representative elements
c. noble gases
b. transition elements
d. nonmetallic elements
____ 103. 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.
____ 104. 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.
____ 105. 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
____ 106. What element in the second period has the largest atomic radius?
a. carbon
c. potassium
b. lithium
d. neon
____ 107. 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.
____ 108. 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
____ 109. 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.
____ 110. 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.
____ 111. Which of the following elements has the smallest ionic radius?
a. Li
c. O
b. K
d. S
____ 112. 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
____ 113. 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.
____ 114. 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
____ 115. Which of the following elements has the smallest first ionization energy?
a. sodium
c. potassium
b. calcium
d. magnesium
____ 116. Which of the following elements has the lowest electronegativity?
a. lithium
c. bromine
b. carbon
d. fluorine
____ 117. 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.
____ 118. 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
____ 119. Which of the following decreases with increasing atomic number in Group 2A?
a. shielding effect
b. ionic size
c. ionization energy
d. number of electrons
____ 120. 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.
____ 121. What is the electron configuration of the calcium ion?
a. 1s 2s 2p 3s 3p
c. 1s 2s 2p 3s 3p 4s
b. 1s 2s 2p 3s 3p 4s
d. 1s 2s 2p 3s
____ 122. 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
____ 123. 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
____ 124. What is the formula of the ion formed when tin achieves a stable electron configuration?
a. Sn
c. Sn
b. Sn
d. Sn
____ 125. 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
____ 126. 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
____ 127. 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
____ 128. 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
____ 129. 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
____ 130. 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
____ 131. What is the formula unit of sodium nitride?
a. NaN
b. Na N
c. Na N
d. NaN
____ 132. What is the formula unit of aluminum oxide?
a. AlO
b. Al O
c. AlO
d. Al O
____ 133. 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
____ 134. What is the formula for sodium sulfate?
a. NaSO
b. Na SO
c. Na(SO )
d. Na (SO )
____ 135. 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.
____ 136. Which of the following compounds has the formula KNO ?
a. potassium nitrate
c. potassium nitrite
b. potassium nitride
d. potassium nitrogen oxide
____ 137. 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.
____ 138. 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.
____ 139. 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
____ 140. Which of the following elements can form diatomic molecules held together by triple covalent bonds?
a. carbon
c. fluorine
b. oxygen
d. nitrogen
____ 141. Which noble gas has the same electron configuration as the oxygen in a water molecule?
a. helium
c. argon
b. neon
d. xenon
____ 142. Which of the following diatomic molecules is joined by a double covalent bond?
a.
c.
b.
d.
____ 143. A molecule with a single covalent bond is ____.
a. CO
c. CO
b. Cl
d. N
____ 144. 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
____ 145. 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.
____ 146. 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
____ 147. Which of the following bonds is the least reactive?
a. C—C
c. O—H
b. H—H
d. H—Cl
____ 148. In which of the following compounds is the octet expanded to include 12 electrons?
a. H S
c. PCl
b. PCl
d. SF
____ 149. 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
____ 150. 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
____ 151. 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
____ 152. Sigma bonds are formed as a result of the overlapping of which type(s) of atomic orbital(s)?
a. s only
b. p only
c. d only
d. s and p
____ 153. 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
____ 154. 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
____ 155. What type of hybrid orbital exists in the methane molecule?
a. sp
c. sp
b. sp
d. sp d
____ 156. What is the shape of a molecule with a triple bond?
a. tetrahedral
c. bent
b. pyramidal
d. linear
____ 157. 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.
____ 158. How many pi bonds are formed when sp hybridization occurs in ethene, C H ?
a. 0
c. 2
b. 1
d. 3
____ 159. Which of the following atoms acquires the most negative charge in a covalent bond with hydrogen?
a. C
c. O
b. Na
d. S
____ 160. 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
____ 161. 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
____ 162. 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.
____ 163. 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
____ 164. 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
c. copper, Cu , cuprous ion
b. zinc, Zn , zincate ion
d. sulfur, S , sulfurous ion
____ 165. 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
____ 166. 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
____ 167. What is the Stock name for chromic ion?
a. chromium(I) ion
b. chromium(II) ion
c. chromium(III) ion
d. chromium(IV) ion
____ 168. 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
____ 169. 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
____ 170. 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
____ 171. 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.
____ 172. Which of the following formulas represents an ionic compound?
a. CS
c. N O
b. BaI
d. PCl
____ 173. Which element, when combined with fluorine, would most likely form an ionic compound?
a. lithium
c. phosphorus
b. carbon
d. chlorine
____ 174. 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
____ 175. 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
____ 176. 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
____ 177. 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
____ 178. Which of the following compounds contains the lead(II) ion?
a. PbO
c. Pb2O
b. PbCl4
d. Pb2S
____ 179. 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
____ 180. What is the correct formula for potassium sulfite?
a. KHSO
c. K SO
b. KHSO
d. K SO
____ 181. 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
____ 182. What type of compound is CuSO ?
a. monotomic ionic
b. polyatomic covalent
c. polyatomic ionic
d. binary molecular
____ 183. Sulfur hexafluoride is an example of a ____.
a. monatomic ion
b. polyatomic ion
c. binary compound
d. polyatomic compound
____ 184. 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
____ 185. Which of the following is a binary molecular compound?
a. BeHCO
c. AgI
b. PCl
d. MgS
____ 186. Which of the following formulas represents a molecular compound?
a. ZnO
c. SO
b. Xe
d. BeF
____ 187. 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
____ 188. 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
____ 189. What is the name of H SO ?
a. hyposulfuric acid
b. hydrosulfuric acid
c. sulfuric acid
d. sulfurous acid
____ 190. 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
____ 191. What is the formula for sulfurous acid?
a. H SO
b. H SO
c. H SO
d. H S
____ 192. What is the formula for phosphoric acid?
a. H PO
b. H PO
c. HPO
d. HPO
____ 193. 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
____ 194. Select the correct formula for sulfur hexafluoride.
a. S F
c. F S
b. F SO
d. SF
____ 195. 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
____ 196. 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.
____ 197. Which of the following is the correct name for N O ?
a. nitrous oxide
b. dinitrogen pentoxide
c. nitrogen dioxide
d. nitrate oxide
____ 198. If the spin of one electron in an orbital is clockwise, what is the spin of the other electron in that orbital?
a. zero
c. counterclockwise
b. clockwise
d. both clockwise and counterclockwise
____ 199. What are quanta of light called?
a. charms
b. excitons
c. muons
d. photons
____ 200. Which scientist developed the quantum mechanical model of the atom?
a. Albert Einstein
c. Niels Bohr
b. Erwin Schrodinger
d. Ernest Rutherford
Short Answer
201. Write the electron configuration for chromium.
202. Tin reacts with fluorine to form two different compounds, A and B. Compound A contains 38.5 g of tin for
each 12.3 g of fluorine. Compound B contains 56.5 g of tin for each 36.2 g of fluorine. What is the lowest
whole-number mass ratio of tin that combines with a given mass of fluorine?
Numeric Response
203. How many electrons are in the highest occupied energy level of a neutral chlorine atom?
204. How many electrons are in the highest occupied energy level of a neutral strontium atom?
Essay
205. Describe the different principles that govern the building of an electron configuration.
206. What is the explanation for the discrete lines in atomic emission spectra?
207. What is the quantum mechanical model?
208. Explain what is meant by the Heisenberg uncertainty principle.
209. Explain how to write an ionic formula, given an anion and a cation. As an example, use the phosphate anion.
Write formulas for the compounds produced by the combinations of these ions. Name the compounds for
which you have written formulas.
210. Name the compounds CuBr , SCl , and BaF . Explain the use or omission of the Roman numeral (II) and
the prefix di-.