Chapter 8: Electron Configuration & Chemical Periodicity
1. “Each electron in an atom must have its own unique set of quantum numbers” is a
statement of
A) the aufbau principle. D) the periodic law.
B) the Pauli exclusion principle. E) Heisenberg’s principle.
C) Hund’s rule.
2. Energy states of atoms containing more than one electron arise from nucleus-electron
and electron-electron interactions. Which of the following statements correctly describes these
effects?
A) Larger nuclear charge lowers energy, more electrons in an orbital lowers energy.
B) Larger nuclear charge lowers energy, more electrons in an orbital increases energy.
C) Smaller nuclear charge lowers energy, more electrons in an orbital lowers energy.
D) Smaller nuclear charge lowers energy, more electrons in an orbital increases
energy.
E) None of the above statements is generally correct.
3. Which one of the following statements about orbital energies is incorrect?
A) In the hydrogen atom, the energy of an orbital depends only on the value of the
quantum number n.
B) In many-electron atoms the energy of an orbital depends on both n and l.
C) Inner electrons shield outer electrons more effectively than do electrons in the
same orbital.
D) The splitting of sublevels in many-electron atoms is explained in terms of the
penetration effect.
E) The energy of a given orbital increases as the nuclear charge Z increases.
4. The effective nuclear charge for an atom is less than the actual nuclear charge due to
A) shielding. D) electron-pair repulsion.
B) penetration. E) relativity.
C) paramagnetism.
5. In many-electron atoms, which quantum numbers specify the energy of an electron?
A) n and l B) n and ml C) l and ml D) n and ms E) n, l, and ml
Chapter 8: Electron Configuration and Chemical Periodicity
6. Which of the following sublevels is filled last? 3d, 4s, 4p, 4d, 5s
A) 3d B) 4s C) 4p D) 4d E) 5s
7. “Electrons added to atomic orbitals of the same energy will remain unpaired with
parallel spins until the subshell is more than half-filled” is a statement of
A) the aufbau principle. D) the periodic law.
B) Hund’s rule. E) the singularity rule.
C) the Pauli exclusion principle.
8. Which one of the following statements about atomic structure and quantum numbers is
incorrect?
A) In a given atom, the maximum number of electrons having principal quantum
number n = 3, is 18.
B) The number of orbitals in a given f subshell is 7.
C) For n = 4, the largest possible value of l is 3.
D) For n = 4, the largest possible value of ml is 2.
E) The following set of quantum numbers for a single orbital is not allowed: n = 3, l =
1, ml = –2.
9. In a single atom, what is the maximum number of electrons which can have quantum
number n = 4?
A) 16 B) 18 C) 32 D) 36 E) none of the above
10. Select the correct set of quantum numbers (n, l, ml, ms) for the highest energy electron in
the ground state of potassium, K.
A) 4, 1, –1,
12
B) 4, 1, 0,
12
C) 4, 0, 1,
12
D) 4, 0, 0,
12
E) 4, 1, 1,
12
11. Select the correct set of quantum numbers (n, l, ml, ms) for the highest energy electron in
the ground state of tin, Sn.
A) 5, 2, –1,
12
B) 5, 2, 0,
12
C) 5, 1, 2,
12
D) 5, 1, 0,
12
E) 5, 2, 1,
12
Chapter 8: Electron Configuration and Chemical Periodicity
12. Select the correct set of quantum numbers (n, l, ml, ms) for the first electron removed in
the formation of a cation for strontium, Sr.
A) 5, 1 , 0, –
12
D) 5, 1, 1,
12
B) 5, 1, 0,
12
E) 5, 0, 0, –
12
C) 5, 0, 1,
12
13. Which of the following electron configurations is impossible?
A) 1s22s22p63s23p4 D) 1s22s22p63s33p4
B) 1s22s22p53s13p4 E) 1s12s22p63s23p4
C) 1s22s22p63s23p63d104s2
14. Select the correct electron configuration for sulfur (Z = 16).
A) 1s21p62s22p6 D) 1s22s22p63s23p4
B) 1s22s22p83s23p4 E) 1s22s22p63s23d4
C) 1s22s22p83s23p2
15. Select the correct electron configuration for Cu (Z = 29).
A) [Ar]4s23d9 D) [Ar]4s24d9
B) [Ar]4s13d10 E) [Ar]5s24d9
C) [Ar]4s24p63d3
16. Select the correct electron configuration for Te (Z = 52).
A) [Kr]5s25p64d8 D) [Kr]5s24f14
B) [Kr]5s25d105p4 E) [Kr]5s24d105p4
C) [Kr]5s24d105p6
17. The electronic structure 1s22s22p63s23p64s23d8 refers to the ground state of
A) Kr. B) Ni. C) Fe. D) Pd. E) none of the above.
18. In the ground state of an atom of silver (Ag), how many electrons will there be with the
quantum number l = 1? (The n, ml, and ms quantum numbers may have any appropriate values.)
A) 9 B) 12 C) 18 D) 24 E) 36
Chapter 8: Electron Configuration and Chemical Periodicity
19. An atom of element number 33 (As) is in its ground electronic state. Which one of the
following sets quantum numbers could not apply to any of its electrons?
A) n = 2 l = 1 ml = –1 ms = +
12
B) n = 3 l = 0 ml = 0 ms = –
12
C) n = 3 l = 2 ml = –2 ms = –
12
D) n = 4 l = 0 ml = 0 ms = –
12
E) n = 4 l = 2 ml = 1 ms = +
12
20. How many valence electrons are there in an atom with the electron configuration [noble
gas]ns2(n – 1)d10np3?
A) 2 B) 3 C) 5 D) 10 E) 15
21. How many valence electrons are there in an atom with the electron configuration [noble
gas]ns2(n – 1)d8?
A) 2 B) 6 C) 8 D) 10 E) none of the above
22. Which of the following electron configurations represents the ground state of an
element?
A) [Ne]3s13p1 D) [Ne]3s23p33d1
B) [He]2s12p3 E) [Ne]3s23p3
C) [Ne]3s23p23d1
23. Which of the following sets of elements have the [noble gas]ns2np2 valence electron
configuration?
A) O, S, Se, Te, Po D) C, Si, Ge, Sn, Pb
B) N, P, As, Sb, Bi E) Ti, Zr, Hf
C) F, Cl, Br, I, At
24. In the electron configuration [Ar]4s23d104p4, which are valence electrons?
A) all of the electrons after the [Ar] D) only the 4p4 electrons
B) only the 4s2 electrons E) both the 4s2 and the 4p4 electrons
C) only the 3d10 electrons
Chapter 8: Electron Configuration and Chemical Periodicity
25. Which of the following electron configurations is correct for the excited state of an
element?
A) [He]2s22p5
B) [Ne]3s23p1
C) [Ar]4s14p1
D) [Kr]5s24d7
E) [He]1p1
26. Which of the following electron configurations is not possible for an atom in an excited
state? A) 1s22s22p63s23p63d104s14p1 D) 1s22s22p63s23p63d104s2
B) 1s22s22p63s13p5 E) 1s22s22p63s23p63d104s14p3
C) 1s22s22p63s23p23d2
27. Which of the following fourth-period elements has the smallest atomic radius?
A) K B) Ti C) Cu D) Ge E) Kr
28. Which of the following elements has the largest atomic size?
A) S B) Ca C) Ba D) Po E) Rn
29. Which of the following elements has the smallest atomic size?
A) Na B) Ar C) K D) Ca E) Kr
30. Which of the following elements has the smallest atomic radius?
A) Li B) Ne C) Rb D) Sr E) Xe
31. Which of the following elements has the greatest atomic radius?
A) Li B) Ne C) Rb D) Sr E) Xe
32. Which one of the following equations correctly represents the process relating to the
ionization energy of X?
A) X(s) → X+(g) + e– D) X–(g) → X(g) + e–
B) X2(g) → X+(g) + X–(g) E) X(g) → X+(g) + e–
C) X(g) + e– → X–(g)
Chapter 8: Electron Configuration and Chemical Periodicity
33. Which of the following elements has the largest first ionization energy?
A) Na B) Cl C) Ca D) Te E) Br
34. Which of the following elements has the smallest first ionization energy?
A) Rb B) Mg C) I D) As E) F
35. When comparing the successive ionization energies of an element, an unusually big
increase in ionization energy is seen when
A) the first valence electron is removed.
B) the second valence electron is removed.
C) the eighth electron of is removed.
D) the first core electron is removed.
E) the last valence electron is removed.
36. Identify the element of Period 2 which has the following successive ionization energies,
in kJ/mol.
IE1, 1314 IE2, 3389 IE3, 5298 IE4, 7471
IE5, 10992 IE6, 13329 IE7, 71345 IE8, 84087
A) Li B) B C) O D) Ne E) none of the above
37. Which of the following elements has the largest second ionization energy (IE2)?
A) Li B) B C) O D) F E) Na
38. Elements with the highest first ionization energies are found in the ___________ region
of the periodic table.
A) lower left B) upper left C) center D) lower right E) upper right
39. Which one of the following equations correctly represents the process involved in the
electron affinity of X?
A) X(g) → X+(g) + e– D) X(g) + e– → X–(g)
B) X+(g) → X+(aq) E) X+(g) + Y–(g) → XY(s)
C) X+(g) + e– → X(g)
40. Select the element with the most negative electron affinity (i.e., accepts an electron most
readily).
A) H B) Li C) C D) F E) Ne
Chapter 8: Electron Configuration and Chemical Periodicity
41. Which of the following has the most negative electron affinity?
A) H B) Li C) Na D) K E) Rb
42. Elements with ________________ first ionization energies and ___________ electron
affinities generally form cations.
A) low, very negative
B) high, positive or slightly negative
C) low, positive or slightly negative
D) high, very negative
E) None of the above is generally correct.
43. Elements with _______________ first ionization energies and ___________ electron
affinities generally form anions.
A) low, very negative
B) high, positive or slightly negative
C) low, positive or slightly negative
D) high, very negative
E) None of the above is generally correct.
44. Select the element with the greatest metallic character.
A) Li B) Ca C) Al D) Pb E) Cs
45. Select the element with the least metallic character.
A) Sn B) Sr C) Tl D) Ge E) Ga
46. Metallic behavior is generally associated with
A) elements with very negative electron affinities.
B) elements with low ionization energies.
C) elements with small atomic radii.
D) elements with unpaired electrons.
E) elements with partially filled p orbitals.
47. The most acidic oxides are formed from elements found in the _________________
region of the periodic table.
A) upper right B) upper left C) center D) lower right E) lower left
Chapter 8: Electron Configuration and Chemical Periodicity
Page 123
48. The most basic oxides are formed from elements found in the __________________
region of the periodic table.
A) upper right B) upper left C) center D) lower right E) lower left
49. Select the most acidic compound from the following:
A) SO2 B) Al2O3 C) CaO D) PbO E) H2O
50. Select the most basic compound from the following:
A) Bi2O3 B) SiO2 C) Cs2O D) Na2O E) H2O
A) Si B) Sr C) Ga D) Cs E) S
52. Which of the following ions will be most likely to form when selenium ionizes?
A) Se6+ B) Se4+ C) Se2+ D) Se2– E) Se4–
53. Which of the following atoms will be diamagnetic?
A) Cr B) Ru C) Fe D) Pt E) Cd
54. Which of the following elements is paramagnetic?
A) Kr B) Zn C) Sr D) V E) Ar
55. Select the paramagnetic ion.
A) Cu+ B) Ag+ C) Fe3+ D) Cd2+ E) Ca2+
56. Select the diamagnetic ion.
A) Cu2+ B) Ni2+ C) Cr3+ D) Sc3+ E) Cr2+
arrangement of relative radii is correct?
A) A2– > B– > C > D+ > E2+ D) A2– < B– < C > D+ > E2+
B) E2+ > D+ > C > B– > A2– E) None of the above is correct.
C) A2– > B– > C < D+ < E2+
Chapter 8: Electron Configuration and Chemical Periodicity
58. What is the correct order of decreasing size of the following ions?
A) P3– > Cl– > K+ > Ca2+ D) K+ > Cl– > P3– > Ca2+
B) Ca2+ > K+ > Cl– > P3– E) None of the above is correct.
C) K+ > Cl– > Ca2+ > P3–
59. Write down the maximum number of electrons in an atom which can have
a. quantum number n = 4.
b. orbital designation 3d.
c. orbital designation 2pz.
60. In a single atom, write down the maximum possible number of electrons
a. with n = 3
b. with the designation 4dxy
c. with n = 5, l = 3, ms = +
12
61. State Hund’s rule, and show how it applies to the ground state of phosphorus atoms.
62. Write down the full electron configuration for ground state atoms of
a. S (element 16).
b. Pd (element 46).
63. Write down the full electron configuration for ground state atoms of
a. element number 12.
b. element number 23.
c. element number 32.
Chapter 8: Electron Configuration and Chemical Periodicity
64. Describe how the atomic radii of main group elements vary with position in the periodic
65. Define what is meant by ionization energy, and write a balanced chemical equation to
represent the relevant process for element X.
66. Describe how the ionization energies of main group elements vary across a period in the
periodic table, using a diagram if necessary.
67. Briefly describe how the atomic radii and ionization energies of group 1A(1) elements
compare with those of group 8A(18). Also, explain why the values of these properties are so
different between these two groups.
68. Define what is meant by electron affinity, and write a balanced chemical equation to
represent the relevant process for element Y.
69. In Mendeleev’s version of the periodic table, the elements were arranged in order of
increasing atomic number.
70. Moseley’s measurements of nuclear charges of the elements provided the basis for
arranging the elements of the periodic table in order of increasing atomic number.
Chapter 8: Electron Configuration and Chemical Periodicity
71. According to the exclusion principle, two is the maximum number of electrons in an
atom which can share the same four quantum numbers.
72. The difference in energies between the 1s and 2s orbitals is due to the penetration effect.
73. In neutral atoms, the 3d orbitals have higher energy than the 4s orbitals.
74. Hund’s rule is used to predict the electron configuration of atoms in excited states.
75. Elements in which the outermost electron has the same principal quantum number n,
show similar chemical properties.
76. The maximum number of electrons in an atom with the same value of n is 2n2.
77. Elements in the same period of the periodic table have similar valence shell electron
configurations.
78. Atomic size decreases across a period due to an increase in the effective nuclear charge,
Zeff.
79. First ionization energies of neutral atoms may be positive or negative.
80. Bigger atoms generally have smaller ionization energies.
81. Electron affinities of neutral atoms may be positive or negative.
82. Electron affinities become increasingly positive toward the right hand end of a period.
83. A negative electron affinity implies that the atom repels an approaching electron.
Chapter 8: Electron Configuration and Chemical Periodicity
84. In moving down a group in the periodic table, the oxides of the elements become more
acidic in nature.
85. In forming ions of the first series of transition metals, the 4s electrons are lost before the
3d electrons.