Deck 5: Atomic Energies and Periodicity

Question 1

Explain the effects of nuclear charge and screening on the energies of electrons.

Answer 1

The higher the value of the quantum number l, the more that orbital is screened by electrons in smaller, more stable orbitals.

Question 2

Understand the relationships between the structure of the periodic table and electron configurations.

Answer 2

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.

Question 3

Use the Pauli exclusion principle, Hund's rule, and the orbital filling order to predict electron configurations of atoms and ions.

Answer 3

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).

Question 4

Relate trends in atomic radius, ionization energy, and electron affinity to nuclear charge and electron configuration.

Question 5

Understand why ionic compounds exist and the energetics of their formation.

Question 6

Understand the trends in atomic radius, ionization energy, and electron affinity and their relationships to nuclear charge and electron configuration.

Question 7

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

Question 8

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.

Question 9

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.

Question 10

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.

Question 11

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.

Question 12

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

Question 13

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.

Question 14

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.

Question 15

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.

Question 16

Which orbital(s) is/are partially filled in chromium atoms?

A) 4s
B) 3p
C) 3s
D) 3d
E) a and d

Question 17

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

Question 18

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

Question 19

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

Question 20

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

Question 21

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

Question 22

Which of the following have 4 valance electrons?

A) Al
B) Si
C) P
D) As
E) Be

Question 23

Which orbitals are partially filled in Zr atoms?

A) 5s
B) 4p
C) 4s
D) 4d
E) a and d

Question 24

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

Question 25

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.

Question 26

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

Question 27

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

Question 28

What is the electron configuration of a silicon atom?

A) 1s²2s²3s²2p⁶3p²
B) 1s²2s²3s²2p⁶3p⁴
C) 1s²2s²2p⁶3s²3p²
D) 1s²2s²2p⁶3p⁴
E) 1s²2s²2p⁶3s⁴

Question 29

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

Question 30

How many unpaired electrons are in a molybdenum atom?

A) 2
B) 3
C) 4
D) 5
E) 6

Question 31

How many unpaired electrons are on a sulphur atom?

A) 2
B) 3
C) 4
D) 5
E) 6

Question 32

How many pairs of valence electrons exist on a Cl^-1 ion?

A) 1
B) 2
C) 3
D) 4
E) 5

Question 33

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

Question 34

How many unpaired electrons are there on a Mn⁺² ion?

A) 3
B) 4
C) 5
D) 6
E) 7

Question 35

How many unpaired electrons are there on a Co⁺² ion?

A) 1
B) 2
C) 3
D) 4
E) 5

Question 36

The two outermost electrons of ground state carbon can be described by quantum numbers (n, l, m_l,m_s)

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)

Question 37

Which of the following is NOT an excited state of Si?

A)
B)
C)
D)
E)

Question 38

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.

Question 39

Which of the following choices has the ions in order of decreasing size?

A) Rb⁺, Br^-, Se^2-
B) Br^-, Se^2-, Rb⁺
C) Br^-, Rb⁺, Se^2-
D) Se^2-, Br^-, Rb⁺
E) Rb⁺, Se^2-, Br^-

Question 40

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

Question 41

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.

Question 42

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.

Question 43

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⁺

Question 44

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

Question 45

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

Question 46

Which of the following has the lowest ionization energy?

A) Mg²⁺
B) Na⁺
C) O^2-
D) Se^2-
E) F^-

Question 47

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

Question 48

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.

Question 49

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.

Question 50

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.

Question 51

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.

Question 52

Which ionic compound will have the largest lattice energy?

A) NaCl
B) LiCl
C) KCl
D) CsCl
E) RbCl

Question 53

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.

Question 54

Which ionic compound will have the smallest lattice energy?

A) KCl
B) CaCl₂
C) FeCl₃
D) MgCl₂
E) NaCl

Question 55

Which ionic compound will have the largest lattice energy?

A) Li₂O
B) MgO
C) FeCl₃
D) MgCl₂
E) Fe₂O₃

Question 56

Which of the following compounds is lime?

A) MgO
B) CaOH
C) MgOH
D) CaO
E) CaCO₃

Question 57

Which of the following compounds is potash?

A) MgO
B) CaO
C) K₂O
D) KOH
E) Ca₅(PO₄)F

Question 58

Carbonates contain which of the following anions?

A) CO₂^-
B) CO₃^-
C) CO₃^2-
D) HCO₃^-
E) HCO₂

Question 59

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

Question 60

Boron has the following photelectron spectrum. Based on this draw what the photoelectron spectrum for carbon should be.

Question 61

What is the definition of a "valence electron"?

Question 62

Draw the ground state energy diagram for the valence electrons of selenium.

Question 63

Draw the ground state energy level diagram for Cu²⁺.

Question 64

Why do the atoms follow the general trend of becoming smaller as they move across a row?

Question 65

Explain why MgN would NOT be expected to be stable.

Question 66

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).

Question 67

What is the major use of sodium carbonate?

Question 68

Why do the s and d block elements readily form cations?

Question 69

Which of the group 15 elements is metallic?

Question 70

Which of the group 14 elements are metalloids?

Question 71

Which groups of the periodic table contain no metals?