The Atom - Multiple Choice Review (No Calculators)
PSI AP Chemistry
1. Which of the following experiments determined the charge/mass ratio of an electron?
a. Gold foil experiment
b. Cathode ray experiments
c. Millikan oil drop experiment
d. Anode ray experiments
2. Which of the following experiments determined the charge of an electron?
a. Gold foil experiment
b. Cathode ray experiments
c. Millikan oil drop experiment
d. Anode ray experiments
3. Rutherford’s experiments on the scattering of a particles by thin metal foils established:
a. The mass and charge of an atom are concentrated in a nucleus.
b. Electrons are fundamental particles of all matter.
c. All electrons have the same charge.
d. Atoms are electrically neutral.
4. Which of the following is an incorrect statement about the Rutherford scattering
experiment?
a. Most of the alpha particles directed at a very thin gold foil passed right through the
foil, undeflected from their straight-line path
b. As a result of the experiment, the atom is pictured as consisting mostly of open
space
c. The few alpha particles that were on a collision course with the gold electrons
were repelled backward at acute angles
d. The experiment led to the conclusion that the nucleus of an atom contains all of the
atom's positive charge
5. Vanadium has two naturally occurring isotopes, 50V with an atomic mass of 49.9472 u
and 51V with an atomic mass of 50.9440 u. The atomic weight of vanadium is 50.9415 u.
The percent abundances of the vanadium isotopes are __________% 50V and __________%
51V.
a. 0.25, 99.75
b. 99.75, 0.25
c. 49, 51
d. 1.0, 99
6. An unknown element is found to have three naturally occurring isotopes with atomic
masses of 35.9675 (0.337%), 37.9627 (0.063%), and 39.9624 (99.600%). Which of the
following is the unknown element?
a. Ar
b. K
c. Cl
d. Ca
7. There are two principal isotopes of indium (atomic weight = 114.82 u). One of these,
113
49In, has an atomic mass of 112.9043 u. The second isotope is most likely to be:
a. 111
49In
b. 112
49In
114
c. 49In
d. 115
49In
Use the mass spectrum of the unknown element below to answer questions 8-11.
Peak
m/z
Intensity
A
B
24
25
100
12.7
C
26
14.1
8. How many stable isotopes are represented on the mass spectrum?
a. 1
b. 2
c. 3
d. 4
9. What is the relative abundance of peak B?
a. 12.7
b. 10.02
c. 11.12
d. 78.86
10. What is the average atomic mass based on the mass spectrum data?
a. 24.32 u
b. 22.99 u
c. 30.84 u
d. 25.00 u
11. To which element does this mass spectrum belong?
a. Na
b. Mg
c. Al
d. Ne
12. Chlorine consists of two isotopes, Cl-35 (34.97 u) and Cl-37 (36.97 u). If the relative
intensity of Cl-35 was 100 on the mass spectrum, what would be the relative intensity of
the Cl-35 peak? Use the average atomic mass of chorine from the periodic table.
a. 30.97
b. 24.00
c. 31.58
d. 35.45
Peak A: 34.97 u, Intensity = 100
Peak B: 36.97 u, Intensity = ?
13. The elements I and Te have similar average atomic masses. A sample that was believed to
be a mixture of I and Te was run through a mass spectrometer, resulting in the following
data. All of the following statements are true. Which would be the best basis for
concluding that the sample was pure Te?
a. Te forms ions with a -2 charge, whereas I forms ions with a -1 charge.
b. Te is more abundant than I in the universe
c. I consists of only one naturally occurring isotope with 74 neutrons, whereas Te
has more than one isotope.
d. I has a higher first ionization energy that Te does
14. What color of visible light has the highest energy?
a. Violet
b. Blue
c. Red
d. Green
15. Of the following transitions in the Bohr hydrogen atom, the __________ transition results in
the emission of photon with the longest wavelength.
a. n = 1 → n = 6
b. n = 1 → n = 4
c. n = 6 → n = 1
d. n = 6 → n = 3
16. The energy (J) required for an electronic transition in a Bohr hydrogen atom from n = 2 to
n = 3 is __________ J.
a. 4.00 × 10-19
b. 3.00 × 10-19
c. -3.00 × 10-19
d. -7.90 × 10-19
17. Calculate the wavelength of the spectral line in the spectrum of hydrogen for which ni = 1
and nf = 3. Will this produce an emission or absorption spectral line?
a. 277 nm, absorption
b. 103 nm, absorption
c. 103 nm, emission
d. 397 nm, absorption
18. Which of the following statements is/are incorrect?
I.
The Bohr model successfully predicts energies in the ultraviolet portion of the
hydrogen spectrum
The Bohr model successfully predicts energies and frequencies of hyperfine
spectral lines
The Bohr model associates electron energy with the radius at which an electron
orbits the nucleus
The Bohr model does not explain the failure of an electron to lose energy as it
travels a circular path around a nucleus
II.
III.
IV.
a.
b.
c.
d.
I. only
II. only
III. only
III. and IV.
Use the following answer choices to complete questions 19-20.
I.
II.
III.
IV.
V.
Atomic mass
Atomic number
Occupied Bohr orbit
Number of electrons in a particular Bohr orbit
Coulombic forces
19. Influences the energy of a PES peak.
a. II only
b. II and IV
c. II, III, and V
d. IV only
20. Influences the intensity of a PES peak.
a. I only
b. II only
c. III only
d. IV only
21. What element is represented by the PES spectrum below?
a. He
b. Li
c. B
d. Be
Use the PES spectrum of Chlorine below to answer questions 22-28 below.
22. _____ Which peak is produced by the orbital furthest from the nucleus?
23. _____ Which peak is produced by the ionization of 5 valence electrons?
24. _____Which peak is produced by the 2s orbital?
25. _____Which peak represents electrons in the lowest energy Bohr orbital?
26. Which of the following correctly describes the ratio of the intensities of peaks C and D?
a. 2:3
c. 3:2
b. 3:1
d. 5:1
27. Which orbital and how many electrons are represented by peak B in the PES spectrum of
chlorine above?
a. 1s, 2 electrons
b. 2s, 2 electrons
c. 3s, 2 electrons
d. 3s, 1 electron
28. The PES spectrum of carbon (C) should have how many distinct peaks?
a. 1
c. 3
b. 2
d. 4
29. The failure of the Bohr model to explain the stability of the Bohr orbits was best resolved
by which hypothesis?
a. Rutherford’s nuclear model
b. Quantum numbers
c. Heisenberg’s Uncertainty Principle
d. de-Broglie’s Wave Equation
Use the following to answer question 30-31.
I.
II.
III.
IV.
The model was based on the wave properties of the electron.
The electron behaved as a particle and can exist only in discrete states
Describes the orientation of electrons around a nucleus
Describes the location of the electron as an orbital in which there is a high probability of
finding an electron
V. Explains the degenerate nature of s, p, d, and f suborbitals
30. Which of the above statements applies to the Quantum model?
a. I and II
b. I, III, IV, and V
c. IV and V
d. V only
31. Which of the above statements applies to the Bohr model?
a. I and II
b. I and III
c. II and III
d. IV and V
Use the following choices to answer questions 32-35 below.
a. 1s22s22p1
b. 1s22p1
c. 1s22s22p63s2
d. 1s22s22p73s1
e. 1s22s22p6
32. _____ Corresponds to a noble gas
33. _____ Represents an impossible configuration
34. _____ Ground state configuration for Mg
35. _____ Represents and atom in an excited state
36. What element has the following electron orbital notation?
a. Cl
b. Be
c. S
d. Se
37. Which one of the following electron configurations for the species in their ground state is
not correct?
a. Ca: 1s22s22p63s23p64s2
b. Bi: [Xe]6s24f145d106p3
c. As: [Ar]4s23d104p3
d. P: 1s22s22p63p5
38. Which of the following is the correct electronic configuration for the iron(III) ion?
a. [Ar] 3d5
b. [Ar] 4s2 3d3
c. [Kr] 4s2 3d3
d. [Kr] 3d5
39. How many valence electrons are present in a Cd atom?
a. 12
b. 10
c. 8
d. 2
Use the PES spectrum of Phosphorus below to answer questions 40-42.
40. Which peak corresponds to the 1s orbital?
a. 1.06
b. 1.95
c. 13.5
d. 18.7
e. 208
41. Which peak corresponds to the valence 3p orbital?
a. 1.06
b. 1.95
c. 13.5
d. 18.7
e. 208
42. Predict the intensity and binding energy for the 3p orbital of sulfur compared to
phosphorus?
a. Higher intensity, higher binding energy
b. Higher intensity, lower binding energy
c. Lower intensity , higher binding energy
d. Lower intensity, lower binding energy
43. Which of the following provides the best explanation for the many chemical similarities
between sodium and potassium?
a. Both have atomic masses between 20 and 40 u
b. Both have an ns1 electron configuration for their highest occupied energy level
c. Both are main group elements
d. The atomic numbers of the two elements differ by less than ten
44. Electrons from which valence shell are removed first when transition metals ionize?
a. s
b. p
c. d
d. f
45. The +3 charge more common than +2 charge for transition metals in the Fe group. What
accounts for this trend?
a. Decreased ionic radius
b. Increased electronegativity
c. Increased stability of “d” orbitals
d. Increased shielding
46. Elements with this valence shell electron configuration will commonly form ions with a -2
charge.
a. ns2
b. ns2np1
c. ns2np2
d. ns2np4
47. An atom with the electron configuration of [X]ns2nd3 would be in the same group as _____
and have a likely charge of ____.
a. Cr, +6
b. Ta, +3
c. Nb, +5
d. V,+4
48. What would be the expected ionic charges for Sn?
a. +1 and +2
b. +1 and +3
c. +2 and +3
d. +2 and +4
49. Which gives the correct orbital and order for the electrons lost in the formation of the
following cations?
a. Ca2+ 4p
b. Pd4+ 5s then 5p
c. Co 3+ 4s then 3d
d. Zn2+ 3d
50. Atoms on the left side of the chart tend to form positive ions because...
a. Their principal energy level is almost empty
b. Their principal energy level is almost full
c. Their atomic number is less than other elements in that period
d. Both B and C
51. Atomic radius generally increases as we move…
a. down a group and from right to left across a period
b. up a group and from left to right across a period
c. down a group and from left to right across a period
d. up a group and from right to left across a period
52. The effective nuclear charge of an atom is primarily affected by
a. inner electrons
b. outer electrons
c. nuclear charge
d. electron distribution
53. Electrons in the 1s subshell are much closer to the nucleus in Ar than in He due to the
larger __________ in Ar.
a. nuclear charge
b. electronegativity
c. azimuthal quantum number
d. electron orbital velocity
54. Of the following, which gives the correct order for atomic radius for Mg, Na, P, Si and Ar?
a. Mg > Na > P > Si > Ar
b. Ar > Si > P > Na > Mg
c. Ar > P > Si > Mg > Na
d. Na > Mg > Si > P > Ar
55. Which of the following is expected to have the largest radius?
a. P3b. S2c. Cld. Ar
e. K+
56. Which of the following is expected to have the smallest radius?
a. S2b. Ca2+
c. Cld. K+
e. P357. Which of the following pairs correctly shows the proper relationship between the two
atoms/ions in terms of atomic/ionic radii?
a. Na < Na+
b. Cl > Clc. Ti < Zn
d. N3- > N
Use the PES spectra below to answer questions 58-60. The spectra represent elements
found in the same period on the periodic table.
58. Which of these two elements would have the smaller atomic radii?
a. Spectrum 1
b. Spectrum 2
c. They are both the same
d. There is not enough information
59. What are the identities of the elements represented by each spectrum respectively?
a. Fluorine and nitrogen
b. Carbon and beryllium
c. Sodium and carbon
d. Fluorine and carbon
60. Which of the following is not true regarding the spectra above?
a. The 2s peak at 1.72 in spectrum 1 has a lower binding energy that the 2s peak at
3.04 in spectrum 2 due to a decreased effective nuclear charge.
b. There are 4 valence electrons represented in spectrum 1
c. The valence electrons in spectrum 1 experience more shielding than those
represented in spectrum 2.
d. The peak at 52.6 in spectrum 2 is held closer to the nucleus that the peak at 28.6 in
spectrum 1.
61. The first five ionization energies, in kJ/mol, for a particular element are shown below.
I1
786
I2
1577
I3
3232
I4
4356
I5
16,091
The element is likely to form an ionic compound in which the charge is:
a. 1+
b. 3+
c. 4+
d. 5+
62. Of the choices below, which gives the order for first ionization energies?
a. Kr > Se > Br > Ga > Ge
b. Kr > Br > Se > Ge > Ga
c. Ga > Ge > Se > Br > Kr
d. Br > Se > Ga > Kr > Ge
63. An element having which of the following electronic configurations would have the
greatest ionization energy?
a. [He] 2s2 2p3
b. [He] 2s2 2p5
c. [Ne] 3s2 3p3
d. [Ne] 3s2 3p5
64. The electron configuration belonging to the atom with the highest second ionization
energy is
a. 1s22s22p63s1
b. 1s22s22p63s2
c. 1s22s22p63s23p1
d. 1s22s22p63s23p5
65. Sodium is much more apt to exist as a cation than is chlorine. This is because __________.
a. chlorine is a gas and sodium is a solid
b. chlorine has a greater ionic radius than sodium does
c. chlorine has a greater atomic radius than sodium
d. chlorine has a greater ionization energy than sodium does
Use the following responses to answer questions 66-68 below.
a.
b.
c.
d.
Cl → Cl+ + eCl + + e- → Cl
Cl + → Cl 2+ + eCl 2+ + e- → Cl +
66. _____ The process representing the first ionization energy of chlorine
67. _____ The process representing the second ionization energy of chlorine
68. _____ The process requiring the most energy
Use the following responses to answer questions 69-71. Use an answer once, more than
once, or not at all.
a.
b.
c.
d.
e.
Mg & Al
As & Se
Cl & F
Cr & Mo
Cu & Ag
69. ____ Show a reversal in the trend for first ionization energy because of electron-electron
repulsions.
70. _____ Show a reversal in the trend for first ionization energy because of shielding by full
orbitals.
71. _____Typically form ions with charges of +3 and +6
Use the following responses to answer questions 72-76. Use an answer once, more than
once, or not at all.
a.
b.
c.
d.
Increased shielding of valence electrons
Effective nuclear charge
½ filled orbitals
Fully filled orbitals
72. _____Why gold is a better conductor of electricity than copper.
73. _____ Why the atomic radius decreases as the atomic number increases from 11-17 across
the periodic table.
74. _____ Why atomic radius increases down a group on the periodic table.
75. _____ Why gallium has an unexpected smaller radius than aluminum?
76. _____ Why magnesium has an unexpectedly higher ionization energy that aluminum.
Use the PES spectrum of Nitrogen and Oxygen below to answer questions 77-79.
77. Which of the following best explains why the 1s peak in the nitrogen spectrum has a
lower binding energy compared to the 1s peak in the oxygen spectrum?
a. Nitrogen atoms have a 3 electrons in their p subshell.
b. Oxygen experience more electron-electron repulsions.
c. Nitrogen has a larger atomic radius.
d. Nitrogen atoms have a smaller effective nuclear charge and Coulombic attraction
than oxygen atoms.
78. Given the fact that both elements are in the same period, what information from the PES
spectra indicates that oxygen would have a smaller atomic radius?
a. The number of peaks
b. The intensity of the peaks
c. Binding energy of the peaks
d. B and C
79. What information from the PES spectra above best accounts for nitrogen commonly
forming 3- ions?
a. The number of peaks
b. The intensity of the core electron peaks
c. The intensity of the valence electron
d. Binding energy of the valence electron peaks
80. What trend would you predict for the binding energies of valence electrons moving down
a group on the periodic table?
a. Increasing binding energy due to increased shielding.
b. Increasing binding energy due to deceased shielding.
c. Decreasing binding energy due to increased shielding.
d. Decreasing binding energy due to decreased nuclear charge.
81. Which of the following elements would require the shortest wavelength of light to ionize?
a. B
b. Al
c. C
d. N
82. What would be the necessary wavelength (nm) required to remove one valence electron
from sodium’s outermost “s” orbital? The first ionization energy for sodium is 8.21x10-19
J.
a. 350 nm
b. 242 nm
c. 124 nm
d. 496 nm
83. Between which two elements is the difference in metallic character the greatest?
a. Rb and O
c. Li and O
b. O and I
d. Li and Rb
84. The list that correctly indicates the order of metallic character is __________.
a. B > N > C
b. F > Cl > S
c. Si > P > S
d. P > S > Se
85. Which of the following best describes why tin considered a metal and carbon a nonmetal
despite being in the same group?
a. Carbon has smaller atomic radium
b. The atomic mass of tin is nearly 10 times that of carbon.
c. Tin has less shielding of its valence electrons and a higher ionization energy.
d. Tin has more shielding of valence electrons and a lower ionization energy.
86. Which of the following orbitals of scandium would require the lowest frequency of light to
ionize?
a. 2p
c. 3p
b. 3s
d. 4s
87. Which of the following best describes the variation of electronegativity with respect to an
element’s position on the periodic table?
a. Increases across a period; increases down a group.
b. Increases across a period; decreases down a group.
c. Decreases across a period; increases down a group.
d. Decreases across a period; decreases down a group.
88. Which of the following groups of elements is arranged correctly in order of increasing
electronegativity?
a. B < O < Al < F
b. Al < B < O < F
c. B < O < F < Al
d. F < O < B < Al
89. An element with a lower electronegativity value will likely have
a. Valence shell PES peaks with low binding energies
b. A lower first ionization energy
c. A low nuclear charge and a high amount of shielding
d. All of the above
Use the following responses to answer for questions 90-92.
a. O
b. Fe
c. Rb
d. Mg
e. N
90. _____What is the most electronegative element of the above?
91. _____Which element readily forms cations with multiple oxidation states?
92. _____Which of the elements above has the smallest ionic radius for its most commonly
found ion?
93. Which of the following BEST explains why silicon has a lower electronegativity than
chlorine?
a. Si has a lower nuclear charge and similar shielding than Cl
b. Si has a higher nuclear charge and similar shielding of Cl
c. Si has the equivalent nuclear charge and less shielding than Cl
d. Si has the equivalent nuclear charge and more shielding than Cl
94. Which is not true of nonmetals?
a. They are poor conductors of heat
b. They are poor conductors of electricity
c. Many are gases at room temperature
d. Most tend to lose electrons readily
95. All of the following statements concerning the characteristics of the halogens are true
EXCEPT:
a. The first ionization energies (potentials) decrease as the atomic numbers of the
halogens increase.
b. They only require one more electron to have full “s” and “p” orbitals.
c. Fluorine atoms have the smallest radii.
d. They readily for cations.
96. Element M reacts with chlorine to form a compound with the formula MCl2. Element M is
more reactive than magnesium and has a smaller radius than barium. This element is ___?
a. Sr
b. K
c. Na
d. Ra
97. The only noble gas that does not have the ns2np6 valence electron configuration is _______.
a. radon
b. neon
c. helium
d. krypton
98. An alkaline earth metal forms a compound with oxygen with the formula __________.
(The symbol M represents any one of the alkaline earth metals.)
a. MO
b. M2O
c. MO2
d. M2O2
99. Alkali metals tend to be more reactive than alkaline earth metals because __________.
a. alkali metals have lower densities
b. alkali metals have lower melting points
c. alkali metals have greater atomic radii
d. alkali metals have lower ionization energies
100.
The element in the periodic table that looks like a metal, is a poor thermal
conductor, and acts as an electrical semiconductor is __________.
a. Sn
b. B
c. As
d. Si
Answer Key
1. B
2. C
3. A
4. C
5. A
6. A
7. D
8. C
9. B
10. A
11. B
12. c
13. C
14. A
15. D
16. B
17. B
18. B
19. C
20. D
21. C
22. A
23. A
24. D
25. E
26. B
27. C
28. C
29. D
30. B
31. C
32. E
33. D
34. C
35. B
36. C
37. D
38. A
39. D
40. E
41. A
42. B
43. B
44. A
45. C
46. D
47. C
48. D
49. C
50. A
51. A
52. A
53. A
54. D
55. A
56. B
57. D
58. B
59. C
60. C
61. C
62. B
63. B
64. A
65. D
66. A
67. C
68. C
69. B
70. A
71. D
72. A
73. B
74. A
75. B
76. D
77. D
78. C
79. C
80. C
81. D
82. B
83. A
84. C
85. D
86. D
87. B
88. B
89. D
90. A
91. B
92. D
93. A
94. D
95. D
96. A
97. C
98. A
99. D
100.
D