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Chemistry A Molecular Approach

Nivaldo J. Tro

Chapter 8

Periodic Properties of the Elements - all with Video Answers

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Chapter Questions

02:29

Problem 1

What are periodic properties?

PS
Praj Sharma
Numerade Educator
01:08

Problem 2

Which periodic property is particularly important to nerve signal transmission? Why?

Keenan Mintz
Keenan Mintz
University of Miami
03:08

Problem 3

Explain the contributions of Johann Dobereiner and John Newlands to the organization of elements according to their properties.

PS
Praj Sharma
Numerade Educator
01:27

Problem 4

Who is credited with arranging the periodic table? How are the elements arranged in the modern periodic table?

Keenan Mintz
Keenan Mintz
University of Miami
03:11

Problem 5

Explain the contributions of Meyer and Moseley to the periodic table.

PS
Praj Sharma
Numerade Educator
02:18

Problem 6

The periodic table is a result of the periodic law. What observations led to the periodic law? What theory explains the underlying reasons for the periodic law?

Keenan Mintz
Keenan Mintz
University of Miami
02:08

Problem 7

What is an electron configuration? Give an example.

Anna Miller
Anna Miller
Numerade Educator
04:31

Problem 8

What is Coulomb's law? Explain how the potential energy of two charged particles depends on the distance between the charged particles and on the magnitude and sign of their charges.

Keenan Mintz
Keenan Mintz
University of Miami
03:14

Problem 9

What is shielding? In an atom, which electrons tend to do the most shielding (core electrons or valence electrons)?

PS
Praj Sharma
Numerade Educator
02:30

Problem 10

What is penetration? How does the penetration of an orbital into the region occupied by core electrons affect the energy of an electron in that orbital?

Keenan Mintz
Keenan Mintz
University of Miami
00:33

Problem 11

Why are the sublevels within a principal level split into different energies for multielectron atoms but not for the hydrogen atom?

Anna Miller
Anna Miller
Numerade Educator
01:46

Problem 12

What is an orbital diagram? Provide an example.

Keenan Mintz
Keenan Mintz
University of Miami
03:30

Problem 13

Why is electron spin important when writing electron configurations? Explain in terms of the Pauli exclusion principle.

PS
Praj Sharma
Numerade Educator
02:21

Problem 14

What are degenerate orbitals? According to Hund's rule, how are degenerate orbitals occupied?

Keenan Mintz
Keenan Mintz
University of Miami
01:54

Problem 15

List all orbitals from 1 s through 5 s according to increasing energy for multielectron atoms.

PS
Praj Sharma
Numerade Educator
01:41

Problem 16

What are valence electrons? Why are they important?

Keenan Mintz
Keenan Mintz
University of Miami
02:58

Problem 17

Copy this blank periodic table onto a sheet of paper and label each of the blocks within the table: $s$ block, $p$ block, $d$ block, and $f$ block.

PS
Praj Sharma
Numerade Educator
02:05

Problem 18

Explain why the $s$ block in the periodic table has only two columns while the $p$ block has six.

Keenan Mintz
Keenan Mintz
University of Miami
04:04

Problem 19

Why do the rows in the periodic table get progressively longer as you move down the table? For example, the first row contains 2 elements, the second and third rows each contain 8 elements,
and the fourth and fifth rows each contain 18 elements. Explain.

Sara Ross
Sara Ross
Numerade Educator
01:49

Problem 20

Describe the relationship between a main-group element's lettered group number (the number of the element's column) and its valence electrons.

Keenan Mintz
Keenan Mintz
University of Miami
03:31

Problem 21

Describe the relationship between an element's row number in the periodic table and the highest principal quantum number in the element's electron configuration. How does this relationship differ for main-group elements, transition elements, and inner transition elements?

PS
Praj Sharma
Numerade Educator
02:23

Problem 22

Which of the transition elements in the first transition series have anomalous electron configurations?

Keenan Mintz
Keenan Mintz
University of Miami
01:21

Problem 23

Describe how to write the electron configuration for an element based on its position in the periodic table.

Anna Miller
Anna Miller
Numerade Educator
01:34

Problem 24

Describe the relationship between the properties of an element and the number of valence electrons that it contains.

Keenan Mintz
Keenan Mintz
University of Miami
03:57

Problem 25

List the number of valence electrons for each family, and explain the relationship between the number of valence electrons and the resulting chemistry of the elements in the family.
\begin{equation}
\begin{array}{ll}{\text { a. alkali metals }} & {\text { b. alkaline earth metals }} \\ {\text { c. halogens }} & {\text { d. oxygen family }}\end{array}
\end{equation}

PS
Praj Sharma
Numerade Educator
02:39

Problem 26

Define atomic radius. For main-group elements, describe the observed trends in atomic radius as you move:
\begin{equation}\begin{array}{l}{\text { a. across a period in the periodic table }} \\ {\text { b. down a column in the periodic table }}\end{array}\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
05:05

Problem 27

What is effective nuclear charge? What is shielding?

PS
Praj Sharma
Numerade Educator
03:35

Problem 28

Use the concepts of effective nuclear charge, shielding, and $n$ value of the valence orbital to explain the trend in atomic radius as you move across a period in the periodic table.

Keenan Mintz
Keenan Mintz
University of Miami
04:05

Problem 29

For transition elements, describe and explain the observed trends in atomic radius as you move:
$\begin{array}{l}{\text { a. across a period in the periodic table }} \\ {\text { b. down a column in the periodic table }}\end{array}$

PS
Praj Sharma
Numerade Educator
02:33

Problem 30

How is the electron configuration of an anion different from that of the corresponding neutral atom? How is the electron configuration of a cation different?

Keenan Mintz
Keenan Mintz
University of Miami
01:04

Problem 31

Describe how to write an electron configuration for a transition metal cation. Is the order of electron removal upon ionization simply the reverse of electron addition upon filling? Why or why not?

Anna Miller
Anna Miller
Numerade Educator
02:01

Problem 32

Describe the relationship between
$\begin{array}{l}{\text { a. the radius of a cation and that of the atom from which forms }} \\ {\text { b. the radius of an anion and that of the atom from which it }} \\ {\text { forms }}\end{array}$

Keenan Mintz
Keenan Mintz
University of Miami
04:50

Problem 33

What is ionization energy? What is the difference between first ionization energy and second ionization energy?

PS
Praj Sharma
Numerade Educator
02:12

Problem 34

What is the general trend in first ionization energy as you move down a column in the periodic table? As you move across a row?

Keenan Mintz
Keenan Mintz
University of Miami
01:22

Problem 35

What are the exceptions to the periodic trends in first ionization energy? Why do they occur?

Anna Miller
Anna Miller
Numerade Educator
02:20

Problem 36

Examination of the first few successive ionization energies for a given element usually reveals a large jump between two ionization energies. For example, the successive ionization energies of
magnesium show a large jump between IE $_{2}$ and $\mathrm{IE}_{3}$ . The successive ionization energies of aluminum show a large jump be tween $\mathrm{IE}_{3}$ and $\mathrm{IE}_{4}$ . Explain why these jumps occur and how you might predict them.

Keenan Mintz
Keenan Mintz
University of Miami
01:10

Problem 37

What is electron affinity? What are the observed periodic trends in electron affinity?

Anna Miller
Anna Miller
Numerade Educator
02:01

Problem 38

What is metallic character? What are the observed periodic trends in metallic character?

Keenan Mintz
Keenan Mintz
University of Miami
05:22

Problem 39

Write a general equation for the reaction of an alkali metal with each substance.
\begin{equation}\begin{array} \\ {\text { a. a halogen } \quad \text { b. water }}\end{array}\end{equation}

PS
Praj Sharma
Numerade Educator
02:13

Problem 40

Write a general equation for the reaction of a halogen with each substance.
\begin{equation}\begin{array}{ll}\\ {\text { a. a metal }} & {\text { b. hydrogen }} & {\text { c. another halogen }}\end{array}\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
02:35

Problem 41

Write the full electron configuration for each element.
a. Si $\quad$ b. $\mathrm{O} \quad$ c. $\mathrm{K} \quad$ d. Ne

PS
Praj Sharma
Numerade Educator
02:49

Problem 42

Write the full electron configuration for each element
\begin{equation}
\begin{array}{l}{\text { a. } C \quad \text { b. } P \quad \text { c. Ar } \quad \text { d. } \mathrm{Na}}\end{array}
\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
02:16

Problem 43

Write the full orbital diagram for each element.
\begin{equation}
\begin{array}{lllll}{\text { a. } N} & {\text { b. } F} & {\text { c. Mg }} & {\text { d. Al }}\end{array}
\end{equation}

Anna Miller
Anna Miller
Numerade Educator
03:57

Problem 44

Write the full orbital diagram for each element.
\begin{equation}
\begin{array}{lllll}{\text { a. } S} & {\text { b. } \mathrm{Ca}} & {\text { c. Ne }} & {\text { d. He }}\end{array}
\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
03:36

Problem 45

Use the periodic table to write an electron configuration for each element. Represent core electrons with the symbol of the previous noble gas in brackets.
\begin{equation}
\begin{array}{lllll}{\text { a. } P} & {\text { b. Ge }} & {\text { c. Zr }} & {\text { d. } 1}\end{array}
\end{equation}

Anna Miller
Anna Miller
Numerade Educator
02:59

Problem 46

Use the periodic table to determine the element corresponding to each electron configuration.
\begin{equation}
\begin{array}{ll}{\text { a. }[\mathrm{Ar}] 4 s^{2} 3 d^{10} 4 p^{6}} & {\text { b. }[\mathrm{Ar}] 4 s^{2} 3 d^{2}} \\ {\text { c. }[\mathrm{Kr}] 5 s^{2} 4 d^{10} 5 p^{2}} & {\text { d. }[\mathrm{Kr}] 5 s^{2}}\end{array}
\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
05:02

Problem 47

Use the periodic table to determine each quantity.
\begin{equation}
\begin{array}{l}{\text { a. the number of } 2 \text { s electrons in Li }} \\ {\text { b. the number of } 3 d \text { electrons in } \mathrm{Cu}}\end{array}
\end{equation}\begin{equation}
\begin{array}{l}{\text { c. the number of } 4 p \text { electrons in Br }} \\ {\text { d. the number of } 4 d \text { electrons in Zr }}\end{array}
\end{equation}

PS
Praj Sharma
Numerade Educator
02:51

Problem 48

Use the periodic table to determine each quantity.
\begin{equation}
\begin{array}{l}{\text { a. the number of } 3 \text { electrons in Mg }} \\ {\text { b. the number of 3d electrons in Cr }}\end{array}
\end{equation}\begin{equation}
\begin{array}{l}{\text { c. the number of } 4 d \text { electrons in } \mathrm{Y}} \\ {\text { d. the number of } 6 p \text { electrons in } \mathrm{Pb}}\end{array}
\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
02:45

Problem 49

Name an element in the fourth period (row) of the periodic table with the following:
\begin{equation}
\begin{array}{l}{\text { a. five valence electrons }} \\ {\text { b. four 4p electrons }}\end{array}
\end{equation}\begin{equation}
\begin{array}{l}{\text { c. three } 3 d \text { electrons }} \\ {\text { d. full s and p sublevels }}\end{array}
\end{equation}

PS
Praj Sharma
Numerade Educator
02:02

Problem 50

Name an element in the third period (row) of the periodic table with the following:
\begin{equation}
\begin{array}{l}{\text { a. three valence electrons }} \\ {\text { b. four } 3 p \text { electrons }}\end{array}
\end{equation}\begin{equation}
\begin{array}{l}{\text { c. six 3p electrons }} \\ {\text { d. two 3s electrons and no 3p electrons }}\end{array}\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
01:47

Problem 51

Determine the number of valence electrons in an atom of each element.
\begin{equation}
\begin{array}{lllll}{\text { a. Ba }} & {\text { b. } \mathrm{Cs}} & {\text { c. } \mathrm{Ni}} & {\text { d. } S}\end{array}\end{equation}

Anna Miller
Anna Miller
Numerade Educator
03:17

Problem 52

Determine the number of valence electrons in an atom of each element. Which elements do you expect to lose electrons in their chemical reactions? Which do you expect to gain electrons?
\begin{equation}\begin{array}{l}\\ {\text { a. Al } \quad \text { b. } \mathrm{Sn} \quad \text { c. Br } \quad \text { d. Se }}\end{array}\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
00:58

Problem 53

Which outer electron configuration would you expect to belong to a reactive metal? To a reactive nonmetal?
\begin{equation}\text { a. }n s^{2} \quad \text { b. } n s^{2} n p^{6} \quad \text { c. } n s^{2} n p^{5} \quad \text { d. } n s^{2} n p^{2}\end{equation}

Anna Miller
Anna Miller
Numerade Educator
01:53

Problem 54

Which outer electron configurations would you expect to belong to a noble gas? To a metalloid?
\begin{equation}
\text { a. }n s^{2} \quad \text { b. } n s^{2} n p^{6} \quad \text { c. } n s^{2} n p^{5} \quad \text { d. } n s^{2} n p^{2}\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
03:57

Problem 55

According to Coulomb's law, which pair of charged particles has the lowest potential energy?
\begin{equation}
\begin{array}{l}{\text { a. a particle with a } 1-\text { charge separated by } 150 \mathrm{pm} \text { from a }} \\ {\text { particle with a } 2+\text { charge }} \\ {\text { b. a particle with a } 1-\text { charge separated by } 150 \mathrm{pm} \text { from a }} \\ {\text { particle with a } 1+\text { charge }} \\ {\text { c. a particle with a } 1-\text { charge separated by } 100 \mathrm{pm} \text { from a }} \\ {\text { particle with a } 3+\text { charge }}\end{array}
\end{equation}

PS
Praj Sharma
Numerade Educator
03:50

Problem 56

According to Coulomb's law, rank the interactions between charged particles from lowest potential energy to highest potential energy.
\begin{equation}
\begin{array}{l}{\text { a. a } 1+\text { charge and a } 1-\text { charge separated by } 100 \mathrm{pm}} \\ {\text { b. } a 2+\text { charge and a } 1-\text { charge separated by } 100 \mathrm{pm}} \\ {\text { c. a } 1+\text { charge and a } 1+\text { charge separated by } 100 \mathrm{pm}} \\ {\text { d. a } 1+\text { charge and a } 1-\text { charge separated by } 200 \mathrm{pm}}\end{array}
\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
00:40

Problem 57

Which experience a greater effective nuclear charge: the valence electrons in beryllium or the valence electrons in nitrogen? Why?

Anna Miller
Anna Miller
Numerade Educator
01:41

Problem 58

Arrange the atoms according to decreasing effective nuclear charge experienced by their valence electrons: $\mathrm{s}, \mathrm{Mg}, \mathrm{Al},$ Si.

Keenan Mintz
Keenan Mintz
University of Miami
05:34

Problem 59

If core electrons completely shielded valence electrons from nuclear charge (i.e., if each core electron reduced nuclear charge by 1 unit) and if valence electrons did not shield one another from nuclear charge at all, what would be the effective nuclear charge experienced by the valence electrons of each atom?
\begin{equation}\begin{array} \\ {\text { a. } \mathrm{K} \quad \text { b. Ca } \quad \text { c. } \mathrm{O} \quad \text { d. } \mathrm{C}}\end{array}\end{equation}

Ronald Prasad
Ronald Prasad
Numerade Educator
02:33

Problem 60

In Section $8.6,$ we estimated the effective nuclear charge on beryllium's valence electrons to be slightly greater than $2+.$ What would a similar process predict for the effective nuclear charge
on boron's valence electrons? Would you expect the effective nuclear charge to be different for boron's 2 s electrons compared to its 2$p$ electron? In what way? (Hint: Consider the shape of the
2$p$ orbital compared to that of the 2 orbital.)

Keenan Mintz
Keenan Mintz
University of Miami
03:12

Problem 61

Choose the larger atom from each pair
\begin{equation}\begin{array} \\ {\text { a. Al or In } \quad \text { b. Si or } \mathrm{N} \quad \text { c. Por Pb } \quad \text { d. Cor } \mathrm{F}}\end{array}\end{equation}

PS
Praj Sharma
Numerade Educator
02:43

Problem 62

Choose the larger atom from each pair, if possible.
\begin{equation}\begin{array} \\ {\text { a. Sn or Si } \text { b. Br or Ga } \quad \text { c. Sn or Bi } \quad \text { d. Se or Sn }}\end{array}\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
01:08

Problem 63

Arrange these elements in order of increasing atomic radius:
\begin{equation}\mathrm{Ca}, \mathrm{Rb}, \mathrm{S}, \mathrm{Si}, \mathrm{Ge}, \mathrm{F}
\end{equation}

Anna Miller
Anna Miller
Numerade Educator
01:27

Problem 64

Arrange these elements in order of decreasing atomic radius:
\begin{equation}\mathrm{Cs}, \mathrm{Sb}, \mathrm{S}, \mathrm{Pb}, \mathrm{Se}\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
02:22

Problem 65

Write the electron configuration for each ion.
\begin{equation}
\begin{array}\\ {\text { a. } \mathrm{O}^{2-} \quad \text { b. Br }^{-} \quad \text { c. Sr }^{2+} \quad \text { d. } \mathrm{Co}^{3+} \quad \text { e. } \mathrm{Cu}^{2+}}\end{array}
\end{equation}

Anna Miller
Anna Miller
Numerade Educator
03:14

Problem 66

Write the electron configuration for each ion.
\begin{equation}\begin{array} \\ {\text { a. } \mathrm{Cl}^{-} \quad \text { b. } \mathrm{P}^{3-} \quad \text { c. } \mathrm{K}^{+} \quad \text { d. } \mathrm{Mo}^{3+} \quad \text { e. } \mathrm{V}^{3+}}\end{array}
\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
02:01

Problem 67

Write orbital diagrams for each ion and indicate whether the ion is diamagnetic or paramagnetic.
\begin{equation}\begin{array}\\ {\text { a. } \mathrm{V}^{5+} \quad \text { b. } \mathrm{Cr}^{3+} \quad \text { c. } \mathrm{Ni}^{2+} \quad \text { d. } \mathrm{Fe}^{3+}}\end{array}\end{equation}

Anna Miller
Anna Miller
Numerade Educator
04:58

Problem 68

Write orbital diagrams for each ion and indicate whether the ion is diamagnetic or paramagnetic.
\begin{equation}\begin{array} \\ {\text { a. } \mathrm{Cd}^{2+} \quad \text { b. } \mathrm{Au}^{+} \quad \text { c. } \mathrm{Mo}^{3+} \quad \text { d. } \mathrm{Zr}^{2+}}\end{array}\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
01:30

Problem 69

Which is the larger species in each pair?
\begin{equation}
\text { a. }\mathrm{Li} \text { or } \mathrm{Li}^{-}\text { b. } \mathrm{I}^{-} \text { or } \mathrm{Cs}^{+} \quad \text { c. } \mathrm{Cror} \mathrm{Cr}^{3+} \quad \text { d. } \mathrm{O} \text { or } \mathrm{O}^{2-}
\end{equation}

Anna Miller
Anna Miller
Numerade Educator
02:34

Problem 70

Which is the larger species in each pair?
\begin{equation}
\begin{array} \\ {\text { a. Sr or } S r^{2+} \quad \text { b. } N \text { or } N^{3-} \quad \text { c. Ni or } N i^{2+} \quad \text { d. } S^{2-} \text { or } C a^{2+}}\end{array}
\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
01:48

Problem 71

Arrange this isoelectronic series in order of decreasing radius:
$\mathrm{F}^{-}, \mathrm{O}^{2-}, \mathrm{Mg}^{2}+, \mathrm{Na}^{+}$

Anna Miller
Anna Miller
Numerade Educator
01:41

Problem 72

Arrange this isoelectronic series in order of increasing atomic radius: $\mathrm{Se}^{2-}, \mathrm{Sr}^{2+}, \mathrm{Rb}^{+}, \mathrm{Br}^{-}$ .

Keenan Mintz
Keenan Mintz
University of Miami
01:23

Problem 73

Choose the element with the higher first ionization energy from each pair.
\begin{equation}
\text{ a. Br or Bi } \quad \text { b. Na or Rb } \quad \text { c. As or At } \quad \text { d. P or Sn}
\end{equation}

Anna Miller
Anna Miller
Numerade Educator
01:51

Problem 74

Choose the element with the higher first ionization energy from each pair.
\begin{equation}
\begin{array} \\ {\text { a. Por I } \quad \text { b. Si or Cl } \text { c. Por Sb } \quad \text { d. Ga or Ge }}\end{array}\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
01:29

Problem 75

Arrange these elements in order of increasing first ionizationenergy: Si, $\mathrm{F}, \mathrm{In}, \mathrm{N}$ .

Anna Miller
Anna Miller
Numerade Educator
01:05

Problem 76

Arrange these elements in order of decreasing first ionization energy: $\mathrm{Cl}, \mathrm{S}, \mathrm{Sn}, \mathrm{Pb} .$

Keenan Mintz
Keenan Mintz
University of Miami
01:34

Problem 77

For each element, predict where the "jump" occurs for successive ionization energies. (For example, does the jump occur between the first and second ionization energies, the second and third, or the third and fourth?
\begin{equation}\begin{array} \\ {\text { a. Be }\quad}{\text { b. N }\quad}{\text { c. O }}\quad{\text { d. Li }}\end{array}\end{equation}

Alexander Cheng
Alexander Cheng
Numerade Educator
01:48

Problem 78

Consider this set of ionization energies.
\begin{equation}
\begin{array}{l}{\mathrm{IE}_{1}=578 \mathrm{kJ} / \mathrm{mol}} \\ {\mathrm{IE}_{2}=1820 \mathrm{kJ} / \mathrm{mol}} \\ {\mathrm{IE}_{3}=2750 \mathrm{kJ} / \mathrm{mol}} \\ {\mathrm{IE}_{4}=11,600 \mathrm{kJ} / \mathrm{mol}}\end{array}
\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
01:18

Problem 79

Choose the element with the more negative (more exothermic) electron affinity from each pair.
\begin{equation}\begin{array} \\ {\text { a. Na or Rb } \quad\text { b. Bors }}\quad\end{array} \text { c. CorN } \quad \text { d. Li or } F\end{equation}

Anna Miller
Anna Miller
Numerade Educator
03:11

Problem 80

Choose the element with the more negative (more exothermic) electron affinity from each pair.
\begin{equation}\text { a. Mg or } \mathrm{S} \quad \text { b. K or } \mathrm{Cs} \quad \text { c. Si or } \mathrm{P}\quad \quad \text { d. Ga or } \mathrm{Br} \quad\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
01:43

Problem 81

Choose the more metallic element from each pair.
\begin{equation}\begin{array} \\ {\text { a. Sr or Sb } \quad \text { b. As or Bi } \quad \text { c. Cl or O }}\end{array}\text { d. S or As }\end{equation}

Anna Miller
Anna Miller
Numerade Educator
02:21

Problem 82

Choose the more metallic element from each pair.
\begin{equation}\begin{array} \\ {\text { a. Sb or Pb } \quad \text { b. Kor Ge } \quad \text { c. Ge or Sb }}\quad \text { d. As or Sn }\end{array}\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
01:45

Problem 83

Arrange these elements in order of increasing metallic character:
Fr, Sb, In, S, Ba, Se.

Anna Miller
Anna Miller
Numerade Educator
01:47

Problem 84

Arrange these elements in order of decreasing metallic character: Sr, $\mathrm{N},$ Si, $\mathrm{P}$ , Ga, Al.

Keenan Mintz
Keenan Mintz
University of Miami
04:12

Problem 85

Write the balanced chemical equation for the reaction of solid strontium with iodine gas.

AS
Alyssa Sherry
Numerade Educator
01:37

Problem 86

Based on the ionization energies of the alkali metals, which alkali metal would you expect to undergo the most exothermic reaction with chlorine gas? Write a balanced chemical equation for the reaction.

Keenan Mintz
Keenan Mintz
University of Miami
01:02

Problem 87

Write the balanced chemical equation for the reaction of solid lithium with liquid water.

Anna Miller
Anna Miller
Numerade Educator
00:56

Problem 88

Write the balanced chemical equation for the reaction of solid potassium with liquid water.

Keenan Mintz
Keenan Mintz
University of Miami
00:50

Problem 89

Write the balanced equation for the reaction of hydrogen gas with bromine gas.

Anna Miller
Anna Miller
Numerade Educator
00:47

Problem 90

Write the balanced equation for the reaction of chlorine gas with fluorine gas.

Keenan Mintz
Keenan Mintz
University of Miami
00:58

Problem 91

Bromine is a highly reactive liquid while krypton is an inert gas. Explain this difference based on their electron configurations.

Anna Miller
Anna Miller
Numerade Educator
01:52

Problem 92

Potassium is a highly reactive metal while argon is an inert gas. Explain this difference based on their electron configurations.

Keenan Mintz
Keenan Mintz
University of Miami
01:32

Problem 93

Both vanadium and its $3+$ ion are paramagnetic. Refer to their electron configurations to explain this statement.

Anna Miller
Anna Miller
Numerade Educator
01:36

Problem 94

Refer to their electron configurations to explain why copper is paramagnetic while its $1+$ ion is not.

Keenan Mintz
Keenan Mintz
University of Miami
01:10

Problem 95

Suppose you were trying to find a substitute for $\mathrm{K}^{+}$ in nerve signal transmission. Where would you begin your search? What ions would be most like $\mathrm{K}^{+}$ For each ion you propose, explain the ways in which it would be similar to $\mathrm{K}^{+}$ and the ways it would be different. Refer to periodic trends in your discussion.

Anna Miller
Anna Miller
Numerade Educator
03:27

Problem 96

Suppose you were trying to find a substitute for $\mathrm{Na}^{+}$ in nerve signal transmission. Where would you begin your search? What ions would be most like Na $^{+} ?$ For each ion you propose, explain the ways in which it would be similar to $\mathrm{Na}^{+}$ and the ways it would be different. Use periodic trends in your discussion.

Keenan Mintz
Keenan Mintz
University of Miami
00:59

Problem 97

Life on Earth evolved based on the element carbon. Based on periodic properties, what two or three elements would you expect to be most like carbon?

Anna Miller
Anna Miller
Numerade Educator
01:46

Problem 98

Which pair of elements would you expect to have the most similar atomic radii, and why?
$\begin{array}{l}{\text { a. Si and Ga }} \\ {\text { b. Si and Ge }} \\ {\text { c. Si and As }}\end{array}$

Keenan Mintz
Keenan Mintz
University of Miami
04:46

Problem 99

Consider these elements: $\mathrm{N}, \mathrm{Mg}, \mathrm{O}, \mathrm{F},$ Al.
$\begin{array}{l}{\text { a. Write the electron configuration for each element. }} \\ {\text { b. Arrange the elements in order of decreasing atomic radius. }} \\ {\text { c. Arrange the elements in order of increasing atomic radius. }} \\ {\text { d. Use the electron configurations in part a to explain the dif- }} \\ {\text { ferences between your answers to parts } \mathrm{b} \text { and } c .}\end{array}$

Ronald Prasad
Ronald Prasad
Numerade Educator
06:45

Problem 100

Consider these elements: $\mathrm{P}, \mathrm{Ca}, \mathrm{Si}, \mathrm{S},$ Ga.
$\begin{array}{l}{\text { a. Write the electron configuration for each element. }} \\ {\text { b. Arrange the elements in order of decreasing atomic radius. }} \\ {\text { c. Arrange the elements in order of increasing anization energy. }} \\ {\text { d. Use the electron configurations in part a to explain the differences}} \\ {\text {between your answers to parts b and c. }}\end{array}$

Keenan Mintz
Keenan Mintz
University of Miami
01:21

Problem 101

Explain why atomic radius decreases as you move to the right across a period for main-group elements but not for transition elements.

Anna Miller
Anna Miller
Numerade Educator
02:53

Problem 102

Explain why vanadium (radius $=134 \mathrm{pm}$ ) and copper (radius =128 $\mathrm{pm}$ ) have nearly identical atomic radii, even though the atomic number of copper is about 25$\%$ higher than that of vanadium. What would you predict about the relative densities of these two metals? Look up the densities in a reference book, periodictable, or on the Internet. Are your predictions correct?

Keenan Mintz
Keenan Mintz
University of Miami
01:15

Problem 103

The lightest noble gases, such as helium and neon, are completely inert - they do not form any chemical compounds whatsoever. The heavier noble gases, in contrast, do form a limited number of compounds. Explain this difference in terms of trends in fundamental periodic properties.

Anna Miller
Anna Miller
Numerade Educator
02:22

Problem 104

The lightest halogen is also the most chemically reactive, and reactivity generally decreases as you move down the column of halogens in the periodic table. Explain this trend in terms of periodic properties.

Keenan Mintz
Keenan Mintz
University of Miami
02:54

Problem 105

Write general outer electron configurations $\left(n s^{x} n p^{y}\right)$ for groups 6 $\mathrm{A}$ and 7 $\mathrm{A}$ in the periodic table. The electron affinity of each group 7 $\mathrm{A}$ element is more negative than that of each corresponding group 6 $\mathrm{A}$ element. Use the electron configurations to explain why this is so.

Ronald Prasad
Ronald Prasad
Numerade Educator
02:42

Problem 106

The electron affinity of each group 5 A element is more positive than that of each corresponding group 4 A element. Use the outer electron configurations for these columns to suggest a reason for this observation.

Keenan Mintz
Keenan Mintz
University of Miami
00:40

Problem 107

The elements with atomic numbers 35 and 53 have similar chemical properties. Based on their electronic configurations, predict the atomic number of a heavier element that also should share these chemical properties.

Anna Miller
Anna Miller
Numerade Educator
03:12

Problem 108

Write the electron configurations of the six cations that form from sulfur by the loss of one to six electrons. For those cations that have unpaired electrons, write orbital diagrams.

Keenan Mintz
Keenan Mintz
University of Miami
04:32

Problem 109

You have cracked a secret code that uses elemental symbols to spell words. The code uses numbers to designate the elemental symbols. Each number is the sum of the atomic number and
the highest principal quantum number of the highest occupied orbital of the element whose symbol is to be used. The message may be written forward or backward. Decode the following messages:
$\begin{array}{l}{\text { a. } 10,12,58,11,7,44,63,66} \\ {\text { b. } 9,99,30,95,19,47,79}\end{array}$

Ronald Prasad
Ronald Prasad
Numerade Educator
02:38

Problem 110

The electron affinity of sodium is lower than that of lithium, while the electron affinity of chlorine is higher than that of fluorine. Suggest an explanation for this observation.

Keenan Mintz
Keenan Mintz
University of Miami
02:24

Problem 111

Use Coulomb's law to calculate the ionization energy in kJ/mol of an atom composed of a proton and an electron separated by 100.00 pm. What wavelength of light has sufficient energy to ionize the atom?

Anna Miller
Anna Miller
Numerade Educator
04:34

Problem 112

The first ionization energy of sodium is 496 $\mathrm{kJ} / \mathrm{mol}$ . Use Coulomb's law to estimate the average distance between the sodium nucleus and the 3 s electron. How does this distance compare to the atomic radius of sodium? Explain the difference.

Keenan Mintz
Keenan Mintz
University of Miami
02:17

Problem 113

Consider the elements: $\mathrm{B}, \mathrm{C}, \mathrm{N}, \mathrm{O}, \mathrm{F}$
$\begin{array}{l}{\text { a. Which element has the highest first ionization energy? }} \\ {\text { b. Which element has the largest atomic radius? }} \\ {\text { c. Which element is most metallic? }} \\ {\text { d. Which element has three unpaired electrons? }}\end{array}$

Anna Miller
Anna Miller
Numerade Educator
03:15

Problem 114

Consider the elements: $\mathrm{Na}, \mathrm{Mg}, \mathrm{Al}, \mathrm{Si}, \mathrm{P}$
$\begin{array}{l}{\text { a. Which element has the highest second ionization energy? }} \\ {\text { b. Which element has the smallest atomic radius? }} \\ {\text { c. Which element is least metallic? }} \\ {\text { d. Which element is diamagnetic? }}\end{array}$

Keenan Mintz
Keenan Mintz
University of Miami
11:35

Problem 115

Consider the densities and atomic radii of the noble gases at $25^{\circ} \mathrm{C} :$
$\begin{array}{l}{\text { a. Estimate the densities of argon and xenon by interpolation }} \\ {\text { from the data. }} \\ {\text { b. Estimate the density of the element with atomic number } 118} \\ {\text { by extrapolation from the data. }}\end{array}$ $\begin{array}{l}{\text { c. Use the molar mass of neon to estimate the mass of a neon }} \\ {\text { atom. Then use the atomic radius of neon to calculate the average}} \\ {\text {density of a neon atom. How does this density compare }} \\ {\text { to the density of neon gas? What does this comparison suggest }} \\ {\text {about the nature of neon gas? }}\end{array}$$\begin{array}{l}{\text { d. Use the densities and molar masses of krypton and neon to }} \\ {\text { calculate the number of atoms of each element found in a }} \\ {\text { volume of } 1.0 \text { L. Use these values to estimate the number of }} \\ {\text { atoms present in } 1.0 \text { L of Ar. Now use the molar mass of argon }} \\ {\text { to estimate the density of Ar. How does this estimate com- }} \\ {\text { pare to that in part a? }}\end{array}$

Ronald Prasad
Ronald Prasad
Numerade Educator
03:56

Problem 116

As you have seen, the periodic table is a result of empirical observation (i.e., the periodic law), but quantum-mechanical theory explains why the table is so arranged. Suppose that, in another universe, quantum theory was such that there were one s orbital but only two p orbitals (instead of three) and only three $d$ orbitals (instead of five). Draw out the first four periods of the periodic table in this alternative universe. Which elements would be the equivalent of the noble gases? Halogens?
Alkali metals?

Keenan Mintz
Keenan Mintz
University of Miami
02:32

Problem 117

Consider the metals in the first transition series. Use periodic trends to predict a trend in density as you move to the right across the series.

Kevin Chimex
Kevin Chimex
Numerade Educator
02:47

Problem 118

Imagine a universe in which the value of $m_{s}$ can be $+\frac{1}{2}, 0,$ and $-\frac{1}{2}$ . Assuming that all the other quantum numbers can take only the values possible in our world and that the Pauli exclusion principle applies, determine:
\begin{equation}
\begin{array}{l}{\text { a. the new electronic configuration of neon }} \\ {\text { b. the atomic number of the element with a completed } n=2} \\ {\text { shell }} \\ {\text { c. the number of unpaired electrons in fluorine }}\end{array}
\end{equation}

Keenan Mintz
Keenan Mintz
University of Miami
02:07

Problem 119

A carbon atom can absorb radiation of various wavelengths with resulting changes in its electron configuration. Write orbital diagrams for the electron configuration of carbon that results from absorption of the three longest wavelengths of radiation it can absorb.

David Collins
David Collins
Numerade Educator
02:04

Problem 120

Only trace amounts of the synthetic element darmstadtium, atomic number 110, have been obtained. The element is so highly unstable that no observations of its properties have been possible. Based on its position in the periodic table, propose three different reasonable valence electron configurations for
this element.

Keenan Mintz
Keenan Mintz
University of Miami
01:40

Problem 121

What is the atomic number of the as yet undiscovered element in which the 8$s$ and 8$p$ electron energy levels fill? Predict the chemical behavior of this element.

Kevin Chimex
Kevin Chimex
Numerade Educator
03:30

Problem 122

The trend in second ionization energy for the elements from lithium to fluorine is not a regular one. Predict which of these elements has the highest second ionization energy and which has the lowest and explain. Of the elements $\mathrm{N}, \mathrm{O},$ and $\mathrm{F},$ O has the highest and $\mathrm{N}$ the lowest second ionization energy. Explain.

Keenan Mintz
Keenan Mintz
University of Miami
05:28

Problem 123

Unlike the elements in groups 1 $\mathrm{A}$ and $2 \mathrm{A},$ those in group 3 $\mathrm{A}$ do
not show a regular decrease in first ionization energy as you move down the column. Explain the irregularities.

Ronald Prasad
Ronald Prasad
Numerade Educator
01:55

Problem 124

Using the data in Figures 8.15 and $8.16,$ calculate $\Delta E$ for the reaction $\mathrm{Na}(g)+\mathrm{Cl}(g) \longrightarrow \mathrm{Na}^{+}(g)+\mathrm{Cl}^{-}(g)$

Keenan Mintz
Keenan Mintz
University of Miami
00:36

Problem 125

Even though adding two electrons to O or S forms an ion with a noble gas electron configuration, the second electron affinity of both of these elements is positive. Explain.

Anna Miller
Anna Miller
Numerade Educator
02:03

Problem 126

In Section 2.7 we discussed the metalloids, which form a diagonal band separating the metals from the nonmetals. There are other instances in which elements such as lithium and magnesium that are diagonal to each other have comparable metallic character. Suggest an explanation for this observation.

Keenan Mintz
Keenan Mintz
University of Miami
00:51

Problem 127

The heaviest known alkaline earth metal is radium, atomic number 88. Find the atomic numbers of the as yet undiscovered next two members of the series.

Anna Miller
Anna Miller
Numerade Educator
01:51

Problem 128

Predict the electronic configurations of the first two excited states (next higher-energy states beyond the ground state) of Pd.

Keenan Mintz
Keenan Mintz
University of Miami
03:15

Problem 129

Table 8.2 does not include francium because none of francium's isotopes are stable. Predict the values of the entries for Fr in Table $8.2 .$ Predict the nature of the products of the reaction of Fr
with: (a) water, (b) oxygen, and (c) chlorine.

Ronald Prasad
Ronald Prasad
Numerade Educator
01:22

Problem 130

From its electronic configuration, predict which of the first 10 elements would be most similar in chemical behavior to the as yet undiscovered element 165.

Keenan Mintz
Keenan Mintz
University of Miami
01:10

Problem 131

Imagine that in another universe atoms and elements are identical to ours, except that atoms with six valence electrons have particular stability (in contrast to our universe where atoms with eight valence electrons have particular stability). Give an example of an element in the alternative universe that corresponds to each of the following:
a. a noble gas b. a reactive nonmetal c. a reactive metal

Anna Miller
Anna Miller
Numerade Educator
03:35

Problem 132

The outermost valence electron in atom A experiences an effective nuclear charge of $2+$ and is on average 225 $\mathrm{pm}$ from the nucleus. The outermost valence electron in atom B experiences an effective nuclear charge of $1+$ and is on average 175 pm from the nucleus. Which atom (A or B) has the higher first ionization energy? Explain.

Keenan Mintz
Keenan Mintz
University of Miami
02:14

Problem 133

Determine whether each statement regarding penetration and shielding is true or false. (Assume that all lower-energy orbitals are fully occupied.)
$\begin{array}{l}{\text { a. An electron in a } 3 \text { orbital is more shielded than an electron }} \\ {\text { in a } 2 \text { s orbital. }} \\ {\text { b. An electron in a } 3 \text { orbital penetrates into the region occupied }} \\ {\text {by core electrons more than electrons in a 3p orbital }} \\ {\text { penetrates into the region occupied by core electrons. }}\end{array}$$
\begin{array}{l}{\text { c. An electron in an orbital that penetrates closer to the nucleus }} \\ {\text { always experiences more shielding than an electron in an orbital }} \\ {\text {that does not penetrate as far. }}\end{array}$$
\begin{array}{l}{\text { d. An electron in an orbital that penetrates close to the nucleus }} \\ {\text { tends to experience a higher effective nuclear charge than an }} \\ {\text { electron in an orbital that does not penetrate close to the }} \\ {\text { nucleus. }}\end{array}$

Ronald Prasad
Ronald Prasad
Numerade Educator
01:38

Problem 134

Give a combination of four quantum numbers that could be assigned to an electron occupying a 5$p$ orbital. Do the same for an electron occupying a 6$d$ orbital.

Keenan Mintz
Keenan Mintz
University of Miami
01:44

Problem 135

Use the trends in ionization energy and electron affinity to explain why calcium fluoride has the formula $\mathrm{CaF}_{2}$ and not $\mathrm{Ca}_{2} \mathrm{F}$ or CaF.

Anna Miller
Anna Miller
Numerade Educator
01:40

Problem 136

In a complete sentence describe the relationship between shielding and penetration.

Keenan Mintz
Keenan Mintz
University of Miami
02:03

Problem 137

Play a game to memorize the order in which orbitals fill. Have each group member in turn state the name of the next orbital to fill and the maximum number of electrons it can hold (for example, e, "1$s$ two," "2$s$ two," "2$p$ six," etc.). If a member gets stuck, other group members can help, consulting Figure 8.5 and the accompanying text summary if necessary. However, when a member gets stuck, the next player starts back at "1$s$ two." Keep going until each group member can list all the orbitals in order up to "6$s$ two."

Ronald Prasad
Ronald Prasad
Numerade Educator
04:56

Problem 138

Sketch a periodic table (without element symbols). Include the correct number of rows and columns in the $s, p, d,$ and $f$ blocks. Shade in the squares for elements that have irregular electron
configurations.

Keenan Mintz
Keenan Mintz
University of Miami
04:27

Problem 139

In complete sentences, explain: (a) why $\mathrm{Se}^{2-}$ and $\mathrm{Br}^{-}$ are about
the same size; (b) why $\mathrm{Br}^{-}$ is slightly smaller than $\mathrm{Se}^{2-} ;$ and $(\mathrm{c})$ which singly charged cation you would expect to be approximately the same size as $\mathrm{Se}^{2-}$ and $\mathrm{Br}^{-}$ and why.

Ronald Prasad
Ronald Prasad
Numerade Educator
03:03

Problem 140

Have each member of your group sketch a periodic table indicating a periodic trend (atomic size, first ionization energy, metallic character, etc.). Have each member present his or her table to the rest of the group and explain the trend based on concepts such as orbital size or effective nuclear charge.

Keenan Mintz
Keenan Mintz
University of Miami
12:32

Problem 141

Figure a $\nabla$ plots the first ionization energies of the elements in period $3 .$ Figure $\mathrm{b}$ > lists some of the binary compounds that form between fluorine and the elements in period $3 .$ Figure $\mathrm{c}$ > plots the electron affinities of the elements in period 3 .
Use the information provided in the figures to answer the following questions:
a. In general, first ionization energy increases as we move from left to right across the third period. Explain.
b. There is a decrease in first ionization from $\mathrm{Mg}$ to $\mathrm{Al}$. There is an increase in electron affinity from $\mathrm{Mg}$ to $\mathrm{Al}$. Explain.
c. The first ionization energy of $\mathrm{S}$ is slightly less than that of $\mathrm{P}$. Explain.
d. After reviewing the information in Figures $\mathrm{a}, \mathrm{b},$ and $\mathrm{c},$ Student X offers the following hypothesis: The number of fluorine atoms that bond to a metal atom is equal to the number of valence electrons of the metal. The number of fluorine atoms that bond to a nonmetal atom is equal to 8 minus the number of valence electrons of the nonmetal.
Student Y offers the following hypothesis: In general, there is an inverse relationship between the electron affinities of the period 3 elements and the number of fluorine atoms that bond to the metal or nonmetal atoms.
Do you agree with Student X's hypothesis, with Student Y's hypothesis, both of them, or neither of them? Explain.
e. Write the formula for a compound that forms between silicon and flourine. Does this compound support Student X's hypothesis, Student Y's hypothesis, or neither hypothesis? Explain.

Ronald Prasad
Ronald Prasad
Numerade Educator