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CHEMISTRY: The Molecular Nature of Matter and Change 2016

Martin S. Silberberg, Patricia G. Amateis

Chapter 8

Electron Configuration and Chemical Periodicity

Educators


Problem 1

What would be your reaction to a claim that a new element had been discovered and it fit between tin (Sn) and antimony (Sb) in the periodic table

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Problem 2

Based on results of his study of atomic x-ray spectra, Moseley discovered a relationship that replaced atomic mass as the criterion for ordering the elements. By what criterion are the elements now ordered in the periodic table? Give an example of a sequence of element order that was confirmed by Moseley’s
findings.

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Problem 3

Before Mendeleev published his periodic table, Dobereiner grouped elements with similar properties into triads, in which the unknown properties of one member could be predicted by averaging known values of the properties of the others. To test this idea, predict the values of the following quantities:
(a) The atomic mass of $\mathrm{K}$ from the atomic masses of Na and Rb
(b) The melting point of $\mathrm{Br}_{2}$ from the melting points of $\mathrm{Cl}_{2}$
$\left(-101.0^{\circ} \mathrm{C}\right)$ and $\mathrm{I}_{2}\left(113.6^{\circ} \mathrm{C}\right)$ (actual value $=-7.2^{\circ} \mathrm{C} )$

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Problem 4

To test Dobereiner's idea (Problem $8.3 ),$ predict:
(a) The boiling point of $\mathrm{HBr}$ from the boiling points of $\mathrm{HCl}$
$\left(-84.9^{\circ} \mathrm{C}\right)$ and $\mathrm{HI}\left(-35.4^{\circ} \mathrm{C}\right)$ (actual value $=-67.0^{\circ} \mathrm{C} )$
(b) The boiling point of $\mathrm{AsH}_{3}$ from the boiling points of $\mathrm{PH}_{3}$
$\left(-87.4^{\circ} \mathrm{C}\right)$ and $\mathrm{SbH}_{3}\left(-17.1^{\circ} \mathrm{C}\right)$ (actual value $=-55^{\circ} \mathrm{C} )$

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Problem 5

Summarize the rules for the allowable values of the four quantum numbers of an electron in an atom.

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Problem 6

Which of the quantum numbers relate(s) to the electron only? Which relate(s) to the orbital?

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Problem 7

State the exclusion principle. What does it imply about the number and spin of electrons in an atomic orbital?

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Problem 8

What is the key distinction between sublevel energies in one-electron species, such as the H atom, and those in manyelectron species, such as the C atom? What factors lead to this distinction? Would you expect the pattern of sublevel energies in Be $^{3+}$ to be more like that in $\mathrm{H}$ or that in $\mathrm{C} ?$ Explain.

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Problem 9

Define shielding and effective nuclear charge. What is the connection between the two

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Problem 10

What is penetration? How is it related to shielding? Use the penetration effect to explain the difference in relative orbital energies of a 3$p$ and a 3$d$ electron in the same atom.

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Problem 11

How many electrons in an atom can have each of the following quantum number or sublevel designations?
(a) $n=2, l=1$
(b) 3$d$
(c) 4$s$

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Problem 12

How many electrons in an atom can have each of the following quantum number or sublevel designations?
(a) $n=2, l=1, m_{l}=0$
(b) 5$p$
(c) $n=4, l=3$

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Problem 13

How many electrons in an atom can have each of the following quantum number or sublevel designations?
(a) 4$p$
(b) $n=3, l=1, m_{l}=+1$
(c) $n=5, l=3$

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Problem 14

How many electrons in an atom can have each of the following quantum number or sublevel designations?
(a) 2$s$
(b) $n=3, l=2$
(c) 6$d$

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Problem 15

State the periodic law, and explain its relation to electron configuration. (Use Na and K in your explanation.)

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Problem 16

State Hund’s rule in your own words, and show its application in the orbital diagram of the nitrogen atom.

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Problem 17

How does the aufbau principle, in connection with the periodic law, lead to the format of the periodic table?

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Problem 18

For main-group elements, are outer electron configurations similar or different within a group? Within a period? Explain

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Problem 19

For which blocks of elements are outer electrons the same as valence electrons? For which are d electrons often included among the valence electrons?

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Problem 20

What is the electron capacity of the nth energy level? What is the capacity of the fourth energy level?

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Problem 21

Write a full set of quantum numbers for the following:
(a) The outermost electron in an Rb atom
(b) The electron gained when an $\mathrm{S}^{-}$ ion becomes an $\mathrm{S}^{2-}$ ion
(c) The electron lost when an Ag atom ionizes
(d) The electron gained when an $\mathrm{F}^{-}$ ion forms from an Fatom

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Problem 22

Write a full set of quantum numbers for the following:
(a) The outermost electron in an Li atom
(b) The electron gained when a Br atom becomes a Br' ion
(c) The electron lost when a Cs atom ionizes
(d) The highest energy electron in the ground-state B atom

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Problem 23

Write the full ground-state electron configuration for each:
$$
\begin{array}{ll}{\text { (a) } \mathrm{Rb}} & {\text { (b) Ge }} & {\text { (c) Ar }}\end{array}
$$

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Problem 24

Write the full ground-state electron configuration for each:
$$
\begin{array}{ll}{\text { (a) } \mathrm{Br}} & {\text { (b) } \mathrm{Mg}} & {\text { (c) Se }}\end{array}
$$

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Problem 25

Write the full ground-state electron configuration for each:
$$
\begin{array}{ll}{\text { (a) } \mathrm{Cl}} & {\text { (b) } \mathrm{Si}} & {\text { (c) Sr }}\end{array}
$$

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Problem 26

Write the full ground-state electron configuration for each:
$$
\begin{array}{ll}{\text { (a) } \mathrm{S}} & {\text { (b) } \mathrm{Kr}} & {\text { (c) } \mathrm{Cs}}\end{array}
$$

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Problem 27

Draw a partial (valence-level) orbital diagram, and write the condensed ground-state electron configuration for each:
$$
\begin{array}{llll}{\text { (a) Ti }} & {\text { (b) } \mathrm{Cl}} & {\text { (c) } \mathrm{V}}\end{array}
$$

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Problem 28

Draw a partial (valence-level) orbital diagram, and write the condensed ground-state electron configuration for each:
$$
\begin{array}{ll}{\text { (a) } \mathrm{Ba}} & {\text { (b) } \mathrm{Co}} & {\text { (c) Ag }}\end{array}
$$

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Problem 29

Draw a partial (valence-level) orbital diagram, and write the condensed ground-state electron configuration for each:
$$
\begin{array}{ll}{\text { (a) } \mathrm{Mn}} & {\text { (b) } \mathrm{P}} & {\text { (c) Fe }}\end{array}
$$

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Problem 30

Draw a partial (valence-level) orbital diagram, and write the condensed ground-state electron configuration for each
$$
\begin{array}{ll}{\text { (a) } \mathrm{Ga}} & {\text { (b) } \mathrm{Zn}} & {\text { (c) } \mathrm{Sc}}\end{array}
$$

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Problem 31

Draw the partial (valence-level) orbital diagram, and write the symbol, group number, and period number of the element:
(a) $[\text { He }] 2 s^{2} 2 p^{4}$
(b) $[\mathrm{Ne}] 3 s^{2} 3 p^{3}$

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Problem 32

Draw the partial (valence-level) orbital diagram, and write the symbol, group number, and period number of the element:
(a) $[\mathrm{Kr}] 5 s^{2} 4 d^{10}$
(b) $[\mathrm{Ar}] 4 s^{2} 3 d^{8}$

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Problem 33

Draw the partial (valence-level) orbital diagram, and write the symbol, group number, and period number of the element:
(a) $[\mathrm{Ne}] 3 s^{2} 3 p^{5}$
(b) $[\text { Ar }] 4 s^{2} 3 d^{10} 4 p^{3}$

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Problem 34

Draw the partial (valence-level) orbital diagram, and write the symbol, group number, and period number of the element:
(a) $[\mathrm{Ar}] 4 s^{2} 3 d^{5}$
(b) $[\mathrm{Kr}] 5 s^{2} 4 d^{2}$

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Problem 35

From each partial (valence-level) orbital diagram, write the condensed electron configuration and group number:
(a)
(b)

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Problem 36

From each partial (valence-level) orbital diagram, write the condensed electron configuration and group number:
(a)
(b)

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Problem 37

How many inner, outer, and valence electrons are present in an atom of each of the following elements?
(a) O
(b) Sn
(c) Ca
(d) Fe
(e) Se

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Problem 38

How many inner, outer, and valence electrons are present in an atom of each of the following elements?
(a) Br
(b) Cs
(c) Cr
(d) Sr
(e) F

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Problem 39

Identify each element below, and give the symbols of the other elements in its group:
(a) $[\text { He }] 2 s^{2} 2 p^{1}$
(b) $[\mathrm{Ne}] 3 s^{2} 3 p^{4}$
(c) $[\mathrm{Xe}] 6 s^{2} 5 d^{1}$

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Problem 40

Identify each element below, and give the symbols of the other elements in its group:
(a) $[\mathrm{Ar}] 4 s^{2} 3 d^{10} 4 p^{4}$
(b) $[\mathrm{Xe}] 6 s^{2} 4 f^{14} 5 d^{2}$
(c) $[\mathrm{Ar}] 4 s^{2} 3 d^{5}$

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Problem 41

Identify each element below, and give the symbols of the other elements in its group
(a) $[\text { He }] 2 s^{2} 2 p^{2}$
(b) $[\text { Ar }] 4 s^{2} 3 d^{3}$
(c) $[\mathrm{Ne}] 3 s^{2} 3 p^{3}$

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Problem 42

Identify each element below, and give the symbols of the other elements in its group:
(a) $[\mathrm{Ar}] 4 s^{2} 3 d^{10} 4 p^{2}$
(b) $[\mathrm{Ar}] 4 s^{2} 3 d^{7}$
(c) $[\mathrm{Kr}] 5 s^{2} 4 d^{5}$

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Problem 43

After an atom in its ground state absorbs energy, it exists in an excited state. Spectral lines are produced when the atom returns to its ground state. The yellow-orange line in the sodium spectrum, for example, is produced by the emission of energy when excited sodium atoms return to their ground state. Write the electron configuration and the orbital diagram of the first excited state of sodium. (Hint: The outermost electron is excited.)

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Problem 44

One reason spectroscopists study excited states is to gain information about the energies of orbitals that are unoccupied in an atom’s ground state. Each of the following electron configurations represents an atom in an excited state. Identify the element, and write its condensed ground-state configuration:
$$
\begin{array}{ll}{\text { (a) } 1 s^{2} 2 s^{2} 2 p^{6} 3 s^{1} 3 p^{1}} & {\text { (b) } 1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{4} 4 s^{1}} \\ {\text { (c) } 1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 4 s^{2} 3 d^{4} 4 p^{1}} & {\text { (d) } 1 s^{2} 2 s^{2} 2 p^{5} 3 s^{1}}\end{array}
$$

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Problem 45

If the exact outer limit of an isolated atom cannot be measured, what criterion can we use to determine atomic radii? What is the difference between a covalent radius and a metallic radius?

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Problem 46

Given the following partial (valence-level) electron configurations, (a) identify each element, (b) rank the four elements in order of increasing atomic size, and (c) rank them in order of increasing ionization energy:

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Problem 47

In what region of the periodic table will you find elements with relatively high IEs? With relatively low IEs?

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Problem 48

(a) Why do successive IEs of a given element always increase? (b) When the difference between successive IEs of a given element is exceptionally large (for example, between $\mathrm{IE}_{1}$ and $\mathrm{IE}_{2}$ of $\mathrm{K}$ ), what do we learn about its electron configuration?
(c) The bars represent the relative magnitudes of the first five ionization energies of an atom:
Identify the element and write its complete electron configuration, assuming it comes from (a) Period 2; (b) Period 3; (c) Period 4

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Problem 49

In a plot of IE1 for the Period 3 elements (see Figure 8.15, p. 340), why do the values for elements in Groups 3A(13) and 6A(16) drop slightly below the generally increasing trend?

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Problem 50

Which group in the periodic table has elements with high (endothermic) $\mathrm{IE}_{1}$ and very negative (exothermic) first electron affinities $\left(\mathrm{EA}_{1}\right) ?$ Give the charge on the ions these atoms form.

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Problem 51

The $\mathrm{EA}_{2}$ of an oxygen atom is positive, even though its $\mathrm{EA}_{1}$ is negative. Why does this change of sign occur? Which other elements exhibit a positive $\mathrm{EA}_{2} ?$ Explain.

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Problem 52

How does d-electron shielding influence atomic size among the Period 4 transition elements?

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Problem 53

Arrange each set in order of increasing atomic size
(a) Rb, K, Cs
(b) C, O, Be
(c) Cl, K, S
(d) Mg, K, Ca

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Problem 54

Arrange each set in order of decreasing atomic size:
(a) Ge, Pb, Sn
(b) Sn, Te, Sr
(c) F, Ne, Na
(d) Be, Mg, Na

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Problem 55

Arrange each set of atoms in order of increasing $\mathrm{IE}_{1} :$
(a) Sr, Ca, Ba
(b) N, B, Ne
(c) Br, Rb, Se
(d) As, Sb, Sn

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Problem 56

Arrange each set of atoms in order of decreasing $\mathrm{IE}_{1} :$
(a) Na, Li, K
(b) Be, F, C
(c) Cl, Ar, Na
(d) Cl, Br, Se

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Problem 57

Write the full electron configuration of the Period 2 element with the following successive IEs (in kJ/mol):
$$
\begin{array}{ll}{\mathrm{IE}_{1}=801} & {\mathrm{IE}_{2}=2427 \quad \mathrm{IE}_{3}=3659} \\ {\mathrm{IE}_{4}=25,022} & {\mathrm{IE}_{5}=32,822}\end{array}
$$

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Problem 58

Write the full electron configuration of the Period 3 element with the following successive IEs (in kJ/mol):
$$
\begin{array}{ll}{\mathrm{IE}_{1}=738} & {\mathrm{IE}_{2}=1450} \\ {\mathrm{IE}_{4}=10,539} & {\mathrm{IE}_{5}=13,628}\end{array} \quad \mathrm{IE}_{3}=7732
$$

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Problem 59

Which element in each of the following sets would you expect to have the highest $\mathrm{IE}_{2} ?$
(a) Na, Mg, Al
(b) Na, K, Fe
(c) Sc, Be, Mg

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Problem 60

Which element in each of the following sets would you expect to have the lowest $\mathrm{IE}_{3} ?$
(a) Na, Mg, Al
(b) K, Ca, Sc
(c) Li, Al, B

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Problem 61

List three ways in which metals and nonmetals differ.

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Problem 62

Summarize the trend in metallic character as a function of position in the periodic table. Is it the same as the trend in atomic size? The trend in ionization energy?

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Problem 63

Explain the relationship between atomic size and reducing strength in Group 1A(1). Explain the relationship between IE and oxidizing strength in Group 7A(17).

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Problem 64

Summarize the acid-base behavior of the main-group metal and nonmetal oxides in water. How does oxide acidity in water change down a group and across a period?

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Problem 65

What ions are possible for the two largest stable elements in Group 4A(14)? How does each arise?

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Problem 66

What is a pseudo–noble gas configuration? Give an example of one ion from Group 3A(13) that has it

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Problem 67

How are measurements of paramagnetism used to support electron configurations derived spectroscopically? Use Cu(I) and Cu(II) chlorides as examples.

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Problem 68

The charges of a set of isoelectronic ions vary from 31 to32. Place the ions in order of increasing size.

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Problem 69

Which element would you expect to be more metallic?
(a) Ca or Rb
(b) Mg or Ra
(c) Br or I

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Problem 70

Which element would you expect to be more metallic?
(a) S or Cl
(b) In or Al
(c) As or Br

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Problem 71

Which element would you expect to be less metallic?
(a) Sb or As
(b) Si or P
(c) Be or Na

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Problem 72

Which element would you expect to be less metallic?
(a) Cs or Rn
(b) Sn or Te
(c) Se or Ge

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Problem 73

Does the reaction of a main-group nonmetal oxide in water produce an acidic or a basic solution? Write a balanced equation for the reaction of a Group 6A(16) nonmetal oxide with water.

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Problem 74

Does the reaction of a main-group metal oxide in water produce an acidic or a basic solution? Write a balanced equation for the reaction of a Group 2A(2) oxide with water.

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Problem 75

Write the charge and full ground-state electron configuration of the monatomic ion most likely to be formed by each:
(a) Cl
(b) N
(c) Br

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Problem 76

Write the charge and full ground-state electron configuration of the monatomic ion most likely to be formed by each:
(a) Rb
(b) N
(c) Br

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Problem 77

Write the charge and full ground-state electron configuration of the monatomic ion most likely to be formed by each:
(a) Al
(b) S
(c) Sr

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Problem 78

Write the charge and full ground-state electron configuration of the monatomic ion most likely to be formed by each.
(a) P
(b) Mg
(c) Se

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Problem 79

How many unpaired electrons are present in a ground-state atom from each of the following groups?
(a) 2A(2)
(b) 5A(15)
(c) 8A(18)
(d) 3A(13)

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Problem 80

How many unpaired electrons are present in a ground-state atom from each of the following groups?
(a) 4A(14)
(b) 7A(17)
(c) 1A(1)
(d) 6A(16)

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Problem 81

Which of these are paramagnetic in their ground state?
(a) Ga
(b) Si
(c) Be
(d) Te

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Problem 82

Are compounds of these ground-state ions paramagnetic?
(a) $\mathrm{Ti}^{2+}$
(b) $Z n^{2+}$
(c) $\mathrm{Ca}^{2+}$
(d) $\operatorname{Sn}^{2+}$

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Problem 83

Write the condensed ground-state electron configurations of these transition metal ions, and state which are paramagnetic:
(a) $\mathrm{V}^{3+}$
(b) $\mathrm{Cd}^{2+}$
(c) $\mathrm{Co}^{3+}$
(d) $\mathrm{Ag}^{+}$

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Problem 84

Write the condensed ground-state electron configurations of these transition metal ions, and state which are paramagnetic:
(a) $\mathrm{M} \mathrm{o}^{3+}$
(b) $\mathrm{Au}^{+}$
(c) $\mathrm{Mn}^{2+}$
(d) $\mathrm{Hf}^{2+}$

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Problem 85

Palladium (Pd; Z 5 46) is diamagnetic. Draw partial orbital diagrams to show which of the following electron configurations is consistent with this fact:
(a) $[\mathrm{Kr}] 5 s^{2} 4 d^{8}$
(b) $[\mathrm{Kr}] 4 d^{10}$
(c) $[\mathrm{Kr}] 5 s^{1} 4 d^{9}$

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Problem 86

Niobium $(\mathrm{Nb} ; Z=41)$ has an anomalous ground-state electron configuration for a Group 5 $\mathrm{B}(5)$ element: $[\mathrm{Kr}] 5 s^{1} 4 d^{4} .$ What is the expected electron configuration for elements in this group? Draw partial orbital diagrams to show how paramagnetic measurements could support niobium's actual configuration.

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Problem 87

Rank the ions in each set in order of increasing size, and explain your ranking:
(a) $\mathrm{Li}^{+}, \mathrm{K}^{+}, \mathrm{Na}^{+}$
(b) $\mathrm{Se}^{2-}, \mathrm{Rb}^{+}, \mathrm{Br}^{-}$
(c) $\mathrm{O}^{2-}, \mathrm{F}^{-}, \mathrm{N}^{3-}$

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Problem 88

Rank the ions in each set in order of decreasing size, and explain your ranking:
(a) $\mathrm{Se}^{2-}, \mathrm{S}^{2-}, \mathrm{O}^{2-}$
(b) $\mathrm{Te}^{2-}, \mathrm{Cs}^{+}, \mathrm{I}^{-}$
(c) $\mathrm{Sr}^{2+}, \mathrm{Ba}^{2+}, \mathrm{Cs}^{+}$

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Problem 89

Name the element described in each of the following
(a) Smallest atomic radius in Group 6 $\mathrm{A}(16)$
(b) Largest atomic radius in Period 6
(c) Smallest metal in Period 3
(d) Highest IE, in Group 4 $\mathrm{A}(14)$
(e) Highest IE_ in Period 5
(f) Most metallic in Group 5 $\mathrm{A}(15)$
(g) Group 3 $\mathrm{A}(13)$ element that forms the most basic oxide
(h) Period 4 element with highest energy level filled
(i) Condensed ground-state electron configuration of $[\mathrm{Ne}] 3 s^{2} 3 p^{2}$
(j) Condensed ground-state electron configuration of $[\mathrm{Kr}] 5 s^{2} 4 d^{6}$
(k) Forms $2+$ ion with electron configuration $\left[\text { Ar } 3 d^{3}\right.$
(l) Period 5 element that forms $3+$ ion with pseudo-noble gas configuration
(m) Period 4 transition element that forms $3+$ diamagnetic ion
(n) Period 4 transition element that forms $2+$ ion with a half-filled $d$ sublevel
(o) Heaviest lanthanide
(p) Period 3 element whose $2-$ ion is isoelectronic with Ar
(q) Alkaline earth metal whose cation is isoelectronic with Kr
(r) Group 5 $\mathrm{A}(15)$ metalloid with the most acidic oxide

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Problem 90

Use electron configurations to account for the stability of the lanthanide ions $\mathrm{Ce}^{4+}$ and $\mathrm{Eu}^{2+}$

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Problem 91

When a nonmetal oxide reacts with water, it forms an oxoacid with the same oxidation number as the nonmetal. Give the name and formula of the oxide used to prepare each of these oxoacids: (a) hypochlorous acid; (b) chlorous acid; (c) chloric acid; (d) perchloric acid; (e) sulfuric acid; (f) sulfurous acid; (g) nitric acid; (h) nitrous acid; (i) carbonic acid; ( j) phosphoric acid.

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Problem 92

A fundamental relationship of electrostatics states that the energy required to separate opposite charges of magnitudes $Q_{1}$ and $Q_{2}$ that are a distance $d$ apart is proportional to $\frac{Q_{1} \times Q_{2}}{d} .$ Use this relationship and any other factors to explain the following:
(a) The $\mathrm{IE}_{2}$ of $\mathrm{He}(Z=2)$ is more than twice the $\mathrm{IE}_{1}$ of $\mathrm{H}(Z=1)$
(b) The $\mathrm{IE}_{1}$ of He is less than twice the $\mathrm{IE}_{1}$ of $\mathrm{H}$

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Problem 93

The energy difference between the 5$d$ and 6$s$ sublevels in gold accounts for its color. Assuming this energy difference is about 2.7 eV (electron volt; $1 \mathrm{eV}=1.602 \times 10^{-19} \mathrm{J} ),$ explain why gold has a warm yellow color.

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Problem 94

Write the formula and name of the compound formed from the following ionic interactions: (a) The $2+$ ion and the $1-$ ion are both isoelectronic with the atoms of a chemically unreactive Period 4 element. $(b)$ The $2+$ ion and the $2-$ ion are both isoelectronic with the Period 3 noble gas. (c) The $2+$ ion is the smallest with a filled $d$ sublevel; the anion forms from the smallest halogen. (d) The ions form from the largest and smallest ionizable atoms in Period 2

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Problem 95

The energy changes for many unusual reactions can be determined using Hess’s law (Section 6.5).
(a) Calculate $\Delta E$ for the conversion of $\mathrm{F}^{-}(g)$ into $\mathrm{F}^{+}(g) .$
(b) Calculate $\Delta E$ for the conversion of $\mathrm{Na}^{+}(g)$ into $\mathrm{Na}^{-}(g) .$

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Problem 96

Discuss each conclusion from a study of redox reactions:
(a) The sulfide ion functions only as a reducing agent.
(b) The sulfate ion functions only as an oxidizing agent.
(c) Sulfur dioxide functions as an oxidizing or a reducing agent.

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Problem 97

The hot glowing gases around the Sun, the corona, can reach millions of degrees Celsius, high enough to remove many electrons from gaseous atoms. Iron ions with charges as high as $14+$ have been observed in the corona. Which ions from Fe't to Fe $^{14+}$ are paramagnetic? Which would be most strongly attracted to a magnetic field?

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Problem 98

There are some exceptions to the trends of first and successive ionization energies. For each of the following pairs, explain which ionization energy would be higher:
(a) $\mathrm{IE}_{1}$ of Ga or $\mathrm{IE}_{1}$ of Ge
(b) $\mathrm{IE}_{2}$ of Ga or $\mathrm{IE}_{2}$ of Ge
(c) $\mathrm{IE}_{3}$ of Ga or $\mathrm{IE}_{3}$ of Ge
(d) $\mathrm{IE}_{4}$ of Ga or $\mathrm{IE}_{4}$ of Ge

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Problem 99

Use Figure 8.16, p. 340, to find: (a) the longest wavelength of electromagnetic (EM) radiation that can ionize an alkali metal atom; (b) the longest wavelength of EM radiation that can ionize an alkaline earth metal atom; (c) the elements, other than the alkali and alkaline earth metals, that could also be ionized by the radiation of part (b); (d) the region of the EM spectrum in which these photons are found.

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Problem 100

Rubidium and bromine atoms are depicted at right. (a) What monatomic ions do they form? (b) What electronic feature characterizes this pair of ions, and which noble gas are they related to? (c) Which
pair best represents the relative ionic sizes?

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Problem 101

Partial (valence-level) electron configurations for four different ions are shown below:
Identify the elements from which the ions are derived, and write the formula of the oxide each ion forms.

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Problem 102

Data from the planet Zog for some main-group elements are shown below (Zoggian units are linearly related to Earth units but are not shown). Radio signals from Zog reveal that balloonium is a monatomic gas with two positive nuclear charges. Use the data to deduce the names that Earthlings give to these elements:

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