Electronic Structure And Periodic – Verified Test Bank | Ch5 - Chemistry Canada 4e | Complete Test Bank by John A. Olmsted. DOCX document preview.
CHAPTER 5
ATOMIC ENERGIES AND PERIODICITY
CHAPTER STUDY OBJECTIVES
1. Explain the effects of nuclear charge and screening on the energies of electrons.
SKILLS TO MASTER: Predicting the effects of screening; drawing electron density plots
KEY CONCEPTS: The higher the value of the quantum number l, the more that orbital is screened by electrons in smaller, more stable orbitals.
2. Understand the relationships between the structure of the periodic table and electron configurations.
SKILLS TO MASTER: Counting valence electrons
KEY CONCEPTS: The Pauli exclusion principle states that each electron in an atom has a unique set of quantum numbers. The Aufbau principle states that electrons are placed into atomic orbitals beginning with the lowest-energy electrons, followed by successively higher-energy electrons. Valence electrons are all those of highest principal quantum number plus those in partially filled d and f orbitals.
3. Use the Pauli exclusion principle, Hund’s rule, and the orbital filling order to predict electron configurations of atoms and ions.
SKILLS TO MASTER: Writing electron configurations; filling orbitals with electrons in increasing order of energy; explaining anomalous electron configurations
KEY CONCEPTS: The most stable configuration involving orbitals of equal energies is the one with the maximum number of electrons with the same spin orientation (Hund’s rule).
4. Relate trends in atomic radius, ionization energy, and electron affinity to nuclear charge and electron configuration.
SKILLS TO MASTER: Predicting periodic variations in atomic radius; predicting periodic variations in ionization energy; predicting periodic variations in electron affinity
KEY CONCEPTS: As the principal quantum number n increases, atomic orbitals become larger and less stable. As the atomic number Z increases, any given atomic orbital becomes smaller and more stable. Atomic size decreases from left to right and increases from top to bottom of the periodic table. First ionization energy increases from left to right across each row and decreases from top to bottom of each column of the periodic table.
5. Understand why ionic compounds exist and the energetics of their formation.
SKILLS TO MASTER: Using a Born–Haber cycle to predict the energy change of a reaction, or of one of the steps of the reaction; explaining the effects of charge and ionic radius on lattice energy
KEY CONCEPTS: The lattice energy is the energy required to separate an ionic compound into its ions. The Born–Haber cycle is an example of the use of Hess’s law.
6. Understand the trends in atomic radius, ionization energy, and electron affinity and their relationships to nuclear charge and electron configuration.
SKILLS TO MASTER: Predicting the charges on ions and knowing some common properties of s-, p-, and d-block elements
Multiple Choice QUESTIONS
1. Consider the atomic orbitals of indium. Which of the following statements are true?
1. The 5p orbital completely screens the 5s orbital.
2. The 5s orbital is more stable than the 5p orbital.
3. The orbitals with n = 4 or less shield the 5s orbital.
4. The 5s orbital is less stable than the 5p orbital.
5. The 5s orbital completely screens the 5p orbital.
a) 1 and 3
b) 3, 4 and 5
c) 2, 3 and 5
d) 2 and 3
e) 1 and 4
Difficulty: Easy
Learning Objective: Explain the effects of nuclear charge and screening on the energies of electrons.
Section Reference: 5.1 Orbital Energies
2. The energies for removal of a 1s electron from a H atom, a He atom and a He+ ion are 2.18x10-18 J, 3.94x10-18 J, and 8.72x10-18 J, respectively. These values indicate
a) ionization energies directly reflect the nuclear charge.
b) partially shielding results from electrons having the same principle and azimuthal quantum numbers.
c) s electrons effectively shield p electrons of the same principal quantum number.
d) shielding results only from electrons having lower principal quantum number.
e) ionization energy is independent of nuclear charge.
Difficulty: Easy
Learning Objective: Explain the effects of nuclear charge and screening on the energies of electrons.
Section Reference: 5.1 Orbital Energies
Feedback: Recognition that a 1s electron is capable of partial shielding of a second 1s electron.
3. The ionization energy for a H atom 1s electron is 2.18x10-18 J and that for a 2p electron is 0.545x10-18J.
a) The 2p electron is shielded by the 1s electron, and as a result IE is lower.
b) A 2p electron is further from the nucleus than a 1s electron would be, and as a result IE is lower.
c) The 2p electron is shielded by the 1s electron and will be more easily removed.
d) The 1s electron is closer to the nucleus than a 2p electron would be, therefore it is easier to remove the 1s electron.
e) Electrons never occupy 2p orbitals in hydrogen.
Difficulty: Easy
Learning Objective: Explain the effects of nuclear charge and screening on the energies of electrons.
Section Reference: 5.1 Orbital Energies
Feedback: a) student should recognize that H has only 1 electron; b) correct; c) student should recognize that H has only 1 electron; d) 1s electron will be harder to remove; e) recognize that excited states exist.
4. The 2s orbital calcium is more stable than the 2p orbital even though the 2p orbital has its maximum electron density closer to the nucleus. The reason for this higher stability is
a) calcium’s valence orbitals have n = 4.
b) the 1s orbital screens the 2p orbital more than the 2s orbital.
c) the 2s orbital has more electron density closer to the nucleus than the 2p orbital.
d) calcium is an exception to usual trends.
e) all s orbitals are more stable than p orbitals.
Difficulty: Easy
Learning Objective: Explain the effects of nuclear charge and screening on the energies of electrons.
Section Reference: 5.1 Orbital Energies
5. The energy required to remove an electron from a helium ion, He+, (8.72 x 10-18 J) is more than twice the energy required to remove an electron from a helium atom (3.94 x 10-18 J). The reason(s) are
a) 1s electrons screen each other.
b) the helium ion has a larger Z.
c) the helium atom has less electron-electron repulsions.
d) the measurement is compromised by the Heisenberg uncertainty principle.
e) helium ions have smaller nuclei.
Difficulty: Easy
Learning Objective: Explain the effects of nuclear charge and screening on the energies of electrons.
Section Reference: 5.1 Orbital Energies
6. Titanium atoms have electrons in 1s, 2s, 2p, 3s, 3p, 4s, and 3d orbitals. Which of the following has the orbitals listed in the order of decreasing screening ability?
a) 1s, 2p, 2s, 3p
b) 2s, 2p, 3d, 3s
c) 2s, 2p, 3s, 3d
d) 1s, 2s, 3p, 3s
e) 3s, 4s, 3p, 3d
Difficulty: Easy
Learning Objective: Explain the effects of nuclear charge and screening on the energies of electrons.
Section Reference: 5.1 Orbital Energies
7. Screening accounts for
a) electron-electron repulsions.
b) different numbers of neutrons and protons in the nucleus.
c) differences in deBroglie wavelength.
d) the nuclear charge of different atoms.
e) the nuclear charge actually experienced by an electron.
Difficulty: Easy
Learning Objective: Explain the effects of nuclear charge and screening on the energies of electrons.
Section Reference: 5.1 Orbital Energies
8. Consider the following statements on screening. Which of the following statements are true?
a) The amount of screening depends only on the size of the orbital.
b) The amount of screening depends only on the shape of the orbital.
c) Electrons occupying orbitals of the same principal quantum number do not screen one another effectively.
d) Electrons with the same azimuthal quantum number but different values of the magnetic quantum number do not screen one another effectively.
e) Electrons in d and f orbitals do not screen electrons in s orbitals.
Difficulty: Medium
Learning Objective: Explain the effects of nuclear charge and screening on the energies of electrons.
Section Reference: 5.1 Orbital Energies
9. The theory that states that no 2 electrons on an atom can have the same set of quantum numbers was written by
a) Hund.
b) Pauli.
c) Avogadro.
d) Aufbau.
e( Heisenberg.
Difficulty: Easy
Learning Objective: Understand the relationships between the structure of the periodic table and electron configurations.
Section Reference: 5.2 Structure of the Periodic Table
10. Which orbital(s) is/are partially filled in chromium atoms?
a) 4s
b) 3p
c) 3s
d) 3d
e) a and d
Difficulty: Medium
Learning Objective: Understand the relationships between the structure of the periodic table and electron configurations.
Section Reference: 5.2 Structure of the Periodic Table
11. What is the maximum number of electrons that can occupy the orbitals with principal quantum number = 4?
a) 2
b) 8
c) 18
d) 32
e) 14
Difficulty: Easy
Learning Objective: Understand the relationships between the structure of the periodic table and electron configurations.
Section Reference: 5.2 Structure of the Periodic Table
12. What is the number of electrons that can occupy an orbital with principal quantum number = 4?
a) 2
b) 8
c) 18
d) 32
e) 14
Difficulty: Easy
Learning Objective: Understand the relationships between the structure of the periodic table and electron configurations.
Section Reference: 5.2 Structure of the Periodic Table
Feedback: Regardless of the type, an orbital can accommodate only 2 electrons.
13. What is the maximum number of electrons that can occupy the orbitals with azimuthal quantum number = 3?
a) 7
b) 14
c) 5
d) 10
e) 6
Difficulty: Easy
Learning Objective: Understand the relationships between the structure of the periodic table and electron configurations.
Section Reference: 5.2 Structure of the Periodic Table
14. Which of the following atoms have six valence electrons?
O, W, Se, Eu, Nd
a) 1 and 3
b) 1, 2, and 3
c) 1, 3 and 4
d) 1, 2, 3 and 5
e) 1, 2, 3, and 4
Difficulty: Medium
Learning Objective: Understand the relationships between the structure of the periodic table and electron configurations.
Section Reference: 5.2 Structure of the Periodic Table
15. Which of the following atoms have three valence electrons?
1. B
2. Li
3. Y
4. Element 104
5. Hf
a) 1 and 4
b) 1, 2, and 3
c) 1, 3 and 5
d) 3 and 5
e) 1 and 5
Difficulty: Medium
Learning Objective: Understand the relationships between the structure of the periodic table and electron configurations.
Section Reference: 5.2 Structure of the Periodic Table
16. Which of the following have 4 valance electrons?
a) Al
b) Si
c) P
d) As
e) Be
Difficulty: Easy
Learning Objective: Understand the relationships between the structure of the periodic table and electron configurations.
Section Reference: 5.2 Structure of the Periodic Table
17. Which orbitals are partially filled in Zr atoms?
a) 5s
b) 4p
c) 4s
d) 4d
e) a and d
Difficulty: Easy
Learning Objective: Understand the relationships between the structure of the periodic table and electron configurations.
Section Reference: 5.2 Structure of the Periodic Table
18. What is the maximum number of electrons that can occupy the orbitals with principal quantum number = 3?
a) 2
b) 8
c) 18
d) 6
e) 10
Difficulty: Easy
Learning Objective: Understand the relationships between the structure of the periodic table and electron configurations.
Section Reference: 5.2 Structure of the Periodic Table
19. The electron configuration [Ar]4s13d5 is
a) an excited state configuration of Cr.
b) the ground state configuration of Cr.
c) the ground state electron configuration of Mn.
d) an excited state electron configuration of Mn.
e) the ground state electron configuration of Mo.
Difficulty: Medium
Learning Objective: Use the Pauli exclusion principle, Hund’s rule, and the orbital filling order to predict electron configurations of atoms and ions.
Section Reference: 5.3 Electron Configurations
Feedback: Must recognize Cr as one of the anomalies.
20. Which of the following is the electron configuration of ground state As atoms?
a) [Ar]4s3d104p4
b) [Ar]3s23d103p3
c) [Ar]4s23d104p3
d) [Ar]4s23d104p4
e) [Ar]4s24p63d7
Difficulty: Easy
Learning Objective: Use the Pauli exclusion principle, Hund’s rule, and the orbital filling order to predict electron configurations of atoms and ions.
Section Reference: 5.3 Electron Configurations
21. Which of the following is the electron configuration of ground state Nb atoms?
a) [Kr]5s24d4
b) [Kr]5s14d2
c) [Kr]5s14d3
d) [Kr]5s24d1
e) [Kr]5s24d3
Difficulty: Easy
Learning Objective: Use the Pauli exclusion principle, Hund’s rule, and the orbital filling order to predict electron configurations of atoms and ions.
Section Reference: 5.3 Electron Configurations
22. What is the electron configuration of a silicon atom?
a) 1s22s23s22p63p2
b) 1s22s23s22p63p4
c) 1s22s22p63s23p2
d) 1s22s22p63p4
e) 1s22s22p63s4
Difficulty: Easy
Learning Objective: Use the Pauli exclusion principle, Hund’s rule, and the orbital filling order to predict electron configurations of atoms and ions.
Section Reference: 5.3 Electron Configurations
23. What is the electron configuration of vanadium using the noble gas configuration?
a) [Ne]4s24d3
b) [Ar]4s23d3
c) [Ar]3d5
d) [Ar]4s24d3
e) [Ar]4s24p3
Difficulty: Easy
Learning Objective: Use the Pauli exclusion principle, Hund’s rule, and the orbital filling order to predict electron configurations of atoms and ions.
Section Reference: 5.3 Electron Configurations
24. How many unpaired electrons are in a molybdenum atom?
a) 2
b) 3
c) 4
d) 5
e) 6
Difficulty: Medium
Learning Objective: Use the Pauli exclusion principle, Hund’s rule, and the orbital filling order to predict electron configurations of atoms and ions.
Section Reference: 5.3 Electron Configurations
25. How many unpaired electrons are on a sulphur atom?
a) 2
b) 3
c) 4
d) 5
e) 6
Difficulty: Easy
Learning Objective: Use the Pauli exclusion principle, Hund’s rule, and the orbital filling order to predict electron configurations of atoms and ions.
Section Reference: 5.3 Electron Configurations
26. How many pairs of valence electrons exist on a Cl-1 ion?
a) 1
b) 2
c) 3
d) 4
e) 5
Difficulty: Easy
Learning Objective: Use the Pauli exclusion principle, Hund’s rule, and the orbital filling order to predict electron configurations of atoms and ions.
Section Reference: 5.3 Electron Configurations
27. Salts of Gd3+ are used in magnetic resonance imaging to enhance the quality of the image. How many unpaired electrons are in a Gd3+ ion?
a) 0
b) 4
c) 6
d) 7
e) 9
Difficulty: Medium
Learning Objective: Use the Pauli exclusion principle, Hund’s rule, and the orbital filling order to predict electron configurations of atoms and ions.
Section Reference: 5.3 Electron Configurations
28. How many unpaired electrons are there on a Mn+2 ion?
a) 3
b) 4
c) 5
d) 6
e) 7
Difficulty: Easy
Learning Objective: Use the Pauli exclusion principle, Hund’s rule, and the orbital filling order to predict electron configurations of atoms and ions.
Section Reference: 5.3 Electron Configurations
29. How many unpaired electrons are there on a Co+2 ion?
a) 1
b) 2
c) 3
d) 4
e) 5
Difficulty: Easy
Learning Objective: Use the Pauli exclusion principle, Hund’s rule, and the orbital filling order to predict electron configurations of atoms and ions.
Section Reference: 5.3 Electron Configurations
30. The two outermost electrons of ground state carbon can be described by quantum numbers (n, l, ml,ms)
a) (2,1,1,+1/2) and (2,1,1,–1/2)
b) (2,1,1,+1/2) and (2,1,0,+1/2)
c) (2,0,0,+1/2) and (2,0,0,–1/2)
d) (2,0,0,+1/2) and (2,1,0,+1/2)
e) (2,1,1,+1/2) and (2,1,0,–1/2)
Difficulty: Medium
Learning Objective: Use the Pauli exclusion principle, Hund’s rule, and the orbital filling order to predict electron configurations of atoms and ions.
Section Reference: 5.3 Electron Configurations
Feedback: Must apply concepts from sections 5.2 and 5.3 to correctly determine the solution to this problem.
31. Which of the following is NOT an excited state of Si?
a)
b)
c)
d)
e)
Difficulty: Medium
Learning Objective: Use the Pauli exclusion principle, Hund’s rule, and the orbital filling order to predict electron configurations of atoms and ions.
Section Reference: 5.3 Electron Configurations
32. Sulphur is smaller than phosphorus because
a) it has fewer unpaired electrons.
b) sulphur’s electrons screen more effectively.
c) the effective nuclear charge of sulphur is larger.
d) the two elements are in the same row.
e) it is in the same period and further to the right in the periodic table.
Difficulty: Easy
Learning Objective: Relate trends in atomic radius, ionization energy, and electron affinity to nuclear charge and electron configuration.
Section Reference: 5.4 Periodicity of Atomic Properties
33. Which of the following choices has the ions in order of decreasing size?
a) Rb+, Br-, Se2-
b) Br-, Se2-, Rb+
c) Br-, Rb+, Se2-
d) Se2-, Br-, Rb+
e) Rb+, Se2-, Br-
Difficulty: Medium
Learning Objective: Relate trends in atomic radius, ionization energy, and electron affinity to nuclear charge and electron configuration.
Section Reference: 5.4 Periodicity of Atomic Properties
34. Which of the following is in the order of increasing ionization energy?
a) I, P, Cl
b) I, Cl, P
c) P, Cl, I
d) Cl, P, I
e) P, I, Cl
Difficulty: Easy
Learning Objective: Relate trends in atomic radius, ionization energy, and electron affinity to nuclear charge and electron configuration.
Section Reference: 5.4 Periodicity of Atomic Properties
35. Ionization energy decreases going down a family (column) because
a) valence orbitals become more stable.
b) screening becomes more effective.
c) orbitals are larger.
d) electron affinities are smaller.
e) number of protons in the nucleus increase.
Difficulty: Easy
Learning Objective: Relate trends in atomic radius, ionization energy, and electron affinity to nuclear charge and electron configuration.
Section Reference: 5.4 Periodicity of Atomic Properties
36. The main reason for the anomalous positive electron affinity of nitrogen atoms is
a) electron-electron repulsion in the nitrogen atom.
b) electron-electron repulsion in the nitrogen ion, N-.
c) a lower than expected screening in the N atom.
d) increased negative ionic charge.
e) electrons are strongly attracted to the nucleus.
Difficulty: Medium
Learning Objective: Relate trends in atomic radius, ionization energy, and electron affinity to nuclear charge and electron configuration.
Section Reference: 5.4 Periodicity of Atomic Properties
37. Which of the following is expected to have the highest ionization energy for the next electron?
a) F
b) Sc2+
c) Ca2+
d) Al3+
e) Ga3+
Difficulty: Medium
Learning Objective: Relate trends in atomic radius, ionization energy, and electron affinity to nuclear charge and electron configuration.
Section Reference: 5.4 Periodicity of Atomic Properties
38. Which answer has the elements arranged in order of least to most negative electron affinity?
a) Br, Cl, F
b) N, O, F
c) Na, Mg, F
d) C, N, F
e) F, N, O
Difficulty: Medium
Learning Objective: Relate trends in atomic radius, ionization energy, and electron affinity to nuclear charge and electron configuration.
Section Reference: 5.4 Periodicity of Atomic Properties
Feedback: a) no clear trends down group; b) correct; c) Mg has near 0 EA; d) EA( C) is more negative than that of N; e) order is reversed.
39. Which answer, has the elements arranged in order of most negative to least negative electron affinity?
a) Ge, As, Se
b) As, Ge, Se
c) As, Se, Ge
d) Se, As, Ge
e) Se, Ge, As
Difficulty: Medium
Learning Objective: Relate trends in atomic radius, ionization energy, and electron affinity to nuclear charge and electron configuration.
Section Reference: 5.4 Periodicity of Atomic Properties
40. Which of the following has the lowest ionization energy?
a) Mg2+
b) Na+
c) O2-
d) Se2-
e) F-
Difficulty: Easy
Learning Objective: Relate trends in atomic radius, ionization energy, and electron affinity to nuclear charge and electron configuration.
Section Reference: 5.4 Periodicity of Atomic Properties
41. Which of the following has the greatest magnitude?
a) the first ionization energy of strontium
b) the first electron affinity of fluorine
c) the second ionization energy of magnesium
d) the first ionization energy of oxygen
e) the third ionization energy of magnesium
Difficulty: Easy
Learning Objective: Relate trends in atomic radius, ionization energy, and electron affinity to nuclear charge and electron configuration.
Section Reference: 5.4 Periodicity of Atomic Properties
42. The primary reason that an ionic compound of formula KCl2 does NOT form is
a) the lattice energy is smaller than that of KCl.
b) the electron affinity of Cl is endothermic.
c) the bond energy of Cl2 is prohibitively large.
d) the total ionization energy to reach K+2 is too high.
e) the second electron affinity of Cl is positive.
Difficulty: Medium
Learning Objective: Understand the trends in atomic radius, ionization energy, and electron affinity and their relationships to nuclear charge and electron configuration.
Section Reference: 5.5 Energetics of Ionic Compounds
43. The primary reason that an ionic compound of formula K2Cl does NOT form is
a) the lattice energy is smaller than that of KCl.
b) the first electron affinity of Cl is endothermic.
c) the bond energy of Cl2 is prohibitively large.
d) the total ionization energy to reach K+2 is too high.
e) the second electron affinity of Cl is positive.
Difficulty: Medium
Learning Objective: Understand the trends in atomic radius, ionization energy, and electron affinity and their relationships to nuclear charge and electron configuration.
Section Reference: 5.5 Energetics of Ionic Compounds
44. The lattice energy of CaCl
a) is much larger than as that of KCl.
b) is smaller than that of CaCl2.
c) is larger than that of NaCl.
d) is independent of the size of the cation.
e) is much smaller than that of KCl.
Difficulty: Medium
Learning Objective: Understand the trends in atomic radius, ionization energy, and electron affinity and their relationships to nuclear charge and electron configuration.
Section Reference: 5.5 Energetics of Ionic Compounds
45. Consider the following elements:
Si, Co, Cl, P, Sr, F, Rb. Select the most accurate statement.
a) Rb, S,r and Co will form ions of +2 charge.
b) Al, Si, P, and Cl will form stable anions.
c) Co, Sr, and Rb will form ionic compounds with Cl and F.
d) Sr and Si will form ionic compounds with Co.
e) Rb, Sr, and Co will form ions of +1 charge.
Difficulty: Easy
Learning Objective: Understand the trends in atomic radius, ionization energy, and electron affinity and their relationships to nuclear charge and electron configuration.
Section Reference: 5.5 Energetics of Ionic Compounds
46. Which ionic compound will have the largest lattice energy?
a) NaCl
b) LiCl
c) KCl
d) CsCl
e) RbCl
Difficulty: Easy
Learning Objective: Understand the trends in atomic radius, ionization energy, and electron affinity and their relationships to nuclear charge and electron configuration.
Section Reference: 5.5 Energetics of Ionic Compounds
47. The alkali metals are found in nature as ionic compounds because
a) it takes very little energy to remove an s electron.
b) the electron affinity for these alkali metals is negative.
c) lattice energy is sufficient to overcome the energy required to form the cation.
d) the cations have a noble gas electron configuration.
e) second ionization energies are very large and positive.
Difficulty: Medium
Learning Objective: Understand the trends in atomic radius, ionization energy, and electron affinity and their relationships to nuclear charge and electron configuration.
Section Reference: 5.5 Energetics of Ionic Compounds
48. Which ionic compound will have the smallest lattice energy?
a) KCl
b) CaCl2
c) FeCl3
d) MgCl2
e) NaCl
Difficulty: Easy
Learning Objective: Understand the trends in atomic radius, ionization energy, and electron affinity and their relationships to nuclear charge and electron configuration.
Section Reference: 5.5 Energetics of Ionic Compounds
49. Which ionic compound will have the largest lattice energy?
a) Li2O
b) MgO
c) FeCl3
d) MgCl2
e) Fe2O3
Difficulty: Easy
Learning Objective: Understand the trends in atomic radius, ionization energy, and electron affinity and their relationships to nuclear charge and electron configuration.
Section Reference: 5.5 Energetics of Ionic Compounds
50. Which of the following compounds is lime?
a) MgO
b) CaOH
c) MgOH
d) CaO
e) CaCO3
Difficulty: Easy
Learning Objective: Understand the periodic trends in ionic charge.
Section Reference: 5.6 Ions and Chemical Periodicity
51. Which of the following compounds is potash?
a) MgO
b) CaO
c) K2O
d) KOH
e) Ca5(PO4)F
Difficulty: Easy
Learning Objective: Understand the periodic trends in ionic charge.
Section Reference: 5.6 Ions and Chemical Periodicity
52. Carbonates contain which of the following anions?
a) CO2-
b) CO3-
c) CO32-
d) HCO3-
e) HCO2‑
Difficulty: Easy
Learning Objective: Understand the periodic trends in ionic charge.
Section Reference: 5.6 Ions and Chemical Periodicity
53. Typically the second row of nonmetals, (C and N) make oxoanions with how many oxygen atoms?
a) 1
b) 2
c) 3
d) 4
e) 5
Difficulty: Medium
Learning Objective: Understand the periodic trends in ionic charge.
Section Reference: 5.6 Ions and Chemical Periodicity
ESSAY QUESTIONS
54. Boron has the following photelectron spectrum. Based on this draw what the photoelectron spectrum for carbon should be.
Difficulty: Medium
Learning Objective: Explain the effects of nuclear charge and screening on the energies of electrons.
Section Reference: 5.1 Orbital Energies
55. What is the definition of a “valence electron”?
Difficulty: Medium
Learning Objective: Use the Pauli exclusion principle, Hund’s rule, and the orbital filling order to predict electron configurations of atoms and ions.
Section Reference: 5.3 Electron Configurations
56. Draw the ground state energy diagram for the valence electrons of selenium.
Difficulty: Easy
Learning Objective: Use the Pauli exclusion principle, Hund’s rule, and the orbital filling order to predict electron configurations of atoms and ions.
Section Reference: 5.3 Electron Configurations
57. Draw the ground state energy level diagram for Cu2+.
Difficulty: Easy
Learning Objective: Use the Pauli exclusion principle, Hund’s rule, and the orbital filling order to predict electron configurations of atoms and ions.
Section Reference: 5.3 Electron Configurations
58. Why do the atoms follow the general trend of becoming smaller as they move across a row?
Difficulty: Easy
Learning Objective: Relate trends in atomic radius, ionization energy, and electron affinity to nuclear charge and electron configuration.
Section Reference: 5.4 Periodicity of Atomic Properties
59. Explain why MgN would NOT be expected to be stable.
Difficulty: Medium
Learning Objective: Relate trends in atomic radius, ionization energy, and electron affinity to nuclear charge and electron configuration.
Section Reference: 5.4 Periodicity of Atomic Properties
60. Using a Born–Haber analysis like that shown in the text, estimate the energy released on forming KCl2 from K(s)and Cl2 where K has a +2 charge (the lattice energy for KCl2 is about 2200 kJ/mole; heat of vaporization (K)= 77.1 kJ/mole; IE1(K) =419 kJ/mole; IE2(K)=3051 kJ/mole; Cl2 bond energy = 240 kJ/mole; EA(Cl)= –348.8 kJ).
Difficulty: Hard
Learning Objective: Understand the trends in atomic radius, ionization energy, and electron affinity and their relationships to nuclear charge and electron configuration.
Section Reference: 5.5 Energetics of Ionic Compounds
61. What is the major use of sodium carbonate?
Difficulty: Easy
Learning Objective: Understand the periodic trends in ionic charge.
Section Reference: 5.6 Ions and Chemical Periodicity
62. Why do the s and d block elements readily form cations?
Difficulty: Medium
Learning Objective: Understand the periodic trends in ionic charge.
Section Reference: 5.6 Ions and Chemical Periodicity
63. Which of the group 15 elements is metallic?
Difficulty: Easy
Learning Objective: Understand the periodic trends in ionic charge.
Section Reference: 5.6 Ions and Chemical Periodicity
64. Which of the group 14 elements are metalloids?
Difficulty: Easy
Learning Objective: Understand the periodic trends in ionic charge.
Section Reference: 5.6 Ions and Chemical Periodicity
65. Which groups of the periodic table contain no metals?
Difficulty: Easy
Learning Objective: Understand the periodic trends in ionic charge.
Section Reference: 5.6 Ions and Chemical Periodicity
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