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Chemistry The Central Science In Si Units

Bruce E. Bursten, Catherine Murphy, H. Eugene LeMay

Chapter 7

Periodic Properties of the Elements - all with Video Answers

Educators

Chapter Questions

04:06

Problem 1

As discussed in the text, we can draw an analogy between the attraction of an electron to a nucleus and the act of perceiving light from a light bulb through a frosted glass shade, as shown in the illustration.
Using the simple method of estimating effective nuclear charge, Equation 7.1 , how does the intensity of the light bulb and/or the thickness of the frosting change in the following cases: (a) Moving from boron to carbon? (b) Moving from boron to aluminum?

William Mills
William Mills
Numerade Educator
01:42

Problem 2

Which of these spheres represents $\mathrm{F}$, which represents $\mathrm{Br}$, and which represents $\mathrm{Br}^{-}$ ?

Lottie Adams
Lottie Adams
Numerade Educator
05:06

Problem 3

Consider the $\mathrm{Mg}^{2+}, \mathrm{Cl}^{-}, \mathrm{K}^{+},$ and $\mathrm{Se}^{2-}$ ions. The four spheres below represent these four ions, scaled according to ionic size. (a) Without referring to Figure 7.8 , match each ion to its appropriate sphere. (b) In terms of size, between which of the spheres would you find the (i) $\mathrm{Ca}^{2+}$ and (ii) $\mathrm{S}^{2-}$ ions?

William Mills
William Mills
Numerade Educator
01:12

Problem 4

In the following reaction which sphere represents a metal and which represents a nonmetal?

Lottie Adams
Lottie Adams
Numerade Educator
03:43

Problem 5

Consider the $\mathrm{A}_{2} \mathrm{X}_{4}$, molecule depicted here, where $\mathrm{A}$ and $\mathrm{X}$ are elements. The $\mathrm{A}-\mathrm{A}$ bond length in this molecule is $d_{1}$, and the four $\mathrm{A}-\mathrm{X}$ bond lengths are each $d_{2} .$ (a) In terms of $d_{1}$ and $d_{2}$, how could you define the bonding atomic radii of atoms A and $X ?(\mathbf{b})$ In terms of $d_{1}$ and $d_{2}$, what would you predict for the $X-X$ bond length of an $X_{2}$ molecule?

William Mills
William Mills
Numerade Educator
04:09

Problem 6

Shown below is a qualitative diagram of the atomic orbital energies for an Na atom. The number of orbitals in each subshell is not shown.
(a) Are all of the subshells for $n=1, n=2,$ and $n=3$ shown? If not, what is missing?
(b) The $2 s$ and $2 p$ energy levels are different. Which of the following is the best explanation for why this is the case? (i) The $2 s$ and $2 p$ energy levels have different energies in the hydrogen atom, so of course they will have different energies in the sodium atom, (ii) The energy of the $2 p$ orbital is higher than that of the $2 s$ in all many-electron atoms. (iii) The 2 s level in Na has electrons in it, whereas the $2 p$ does not.
(c) Which of the energy levels holds the highest-energy electron in a sodium atom?
(d) A sodium vapor lamp (Figure 7.23 ) operates by using electricity to excite the highest-energy electron to the next highest-energy level. Light is produced when the excited electron drops back to the lower level. Which two energy levels are involved in this process for the Na atom?

Ronald Prasad
Ronald Prasad
Numerade Educator
02:56

Problem 7

Which of the following charts below shows the general periodic trends for each of the following properties of the main-group elements (you can neglect small deviations going either across a row or down a column of the periodic table)? (1) Bonding atomic radius, (2) first ionization energy, (3) effective nuclear charge.

William Mills
William Mills
Numerade Educator
02:10

Problem 8

An element $\mathrm{X}$ reacts with $\mathrm{F}_{2}(g)$ to form the molecular product shown here. (a) Write a balanced equation for this reaction (do not worry about the phases for $\mathrm{X}$ and the product). (b) Do you think that $\mathrm{X}$ is a metal or nonmetal?

Ronald Prasad
Ronald Prasad
Numerade Educator
03:37

Problem 9

(a) Evaluate the expressions $2 \times 1,2 \times(1+3)$, $2 \times(1+3+5)$, and $2 \times(1+3+5+7)$ (b) How do the atomic numbers of the noble gases relate to the numbers from part (a)? (c) What topic discussed in Chapter 6 is the source of the number "2" in the expressions in part (a)?

William Mills
William Mills
Numerade Educator
01:28

Problem 10

The prefix eka- comes from the Sanskrit word for "one." Mendeleev used this prefix to indicate that the unknown element was one place away from the known element that followed the prefix. For example, eka-silicon, which we now call germanium, is one element below silicon. Mendeleev also predicted the existence of eka-manganese, which was not experimentally confirmed until 1937 because this element is radioactive and does not occur in nature. Based on the periodic table shown in Figure $7.1,$ what do we now call the element Mendeleev called eka-manganese?

James Irizarry
James Irizarry
Numerade Educator
01:35

Problem 11

(a) The five most abundant elements in the Earth's crust are $\mathrm{O}, \mathrm{Si}, \mathrm{Al}, \mathrm{Fe},$ and Ca. Referring to Figure $7.1,$ are any of these elements among those known before $1700 ?$ If so which ones? (b) Seven of the nine elements known since ancient times are metals. Referring to Table $4.5,$ are these metals mostly found at the bottom or top of the activity series?

William Mills
William Mills
Numerade Educator
02:50

Problem 12

Moseley's experiments on $X$ rays emitted from atoms led to the concept of atomic numbers. (a) If arranged in order of increasing atomic mass, which element would come after chlorine? (b) Describe two ways in which the properties of this element differ from the other elements in group $8 \mathrm{~A}$.

James Irizarry
James Irizarry
Numerade Educator
02:26

Problem 13

Among the elements $\mathrm{N}, \mathrm{O}, \mathrm{P},$ and $\mathrm{S},$ which element or elements have the smallest effect nuclear charge if we use Equation 7.1 to calculate $Z_{\text {eff }}$. Which element or elements have the largest effective nuclear charge?

Tracy Tourville
Tracy Tourville
Numerade Educator
12:17

Problem 14

Which of the following statements about effective nuclear charge for the outermost valence electron of an atom is incorrect? (i) The effective nuclear charge can be thought of as the true nuclear charge minus a screening constant due to the other electrons in the atom. (ii) Effective nuclear charge increases (iii) Valence going left to right across a row of the periodic table. electrons screen the nuclear charge more effectively than do core electrons. (iv) The effective nuclear charge shows a sudden decrease when we go from the end of one row to the beginning of the next row of the periodic table. $(\mathbf{v})$ The change in effective nuclear charge going down a column of the periodic table is generally less than that going across a row of the periodic table.

James Irizarry
James Irizarry
Numerade Educator
05:25

Problem 15

Detailed calculations show that the value of $Z_{\text {eff }}$ for the outermost electrons in Na and $\mathrm{K}$ atoms is $2.51+$ and $3.49+$, respectively. (a) What value do you estimate for $Z_{\text {eff }}$ experienced by the outermost electron in both $\mathrm{Na}$ and $\mathrm{K}$ by assuming core electrons contribute 1.00 and valence electrons contribute 0.00 to the screening constant? (b) What values do you estimate for $Z_{\text {eff }}$ using Slater's rules? (c) Which approach gives a more accurate estimate of $Z_{\text {eff }}$ ? (d) Does either method of approximation account for the gradual increase in $Z_{\text {eff }}$ that occurs upon moving down a group? (e) Predict $Z_{\text {eff }}$ for the outermost electrons in the $\mathrm{Rb}$ atom based on the calculations for $\mathrm{Na}$ and $\mathrm{K}$.

William Mills
William Mills
Numerade Educator
14:36

Problem 16

Detailed calculations show that the value of $Z_{\text {eff }}$ for the outermost electrons in $\mathrm{Si}$ and $\mathrm{Cl}$ atoms is $4.29+$ and $6.12+,$ respectively. (a) What value do you estimate for $Z_{\text {eff }}$ experienced by the outermost electron in both Si and Cl by assuming core electrons contribute 1.00 and valence electrons contribute 0.00 to the screening constant? (b) What values do you estimate for $Z_{\text {eff }}$ using Slater's rules? (c) Which approach gives a more accurate estimate of $Z_{\text {eff }}$ ? (d) Which method of approximation more accurately accounts for the steady increase in $Z_{\text {eff }}$ that occurs upon moving left to right across a period? (e) Predict $Z_{\text {eff }}$ for a valence electron in P, phosphorus, based on the calculations for $\mathrm{Si}$ and $\mathrm{Cl}$.

James Irizarry
James Irizarry
Numerade Educator
01:54

Problem 17

Which will experience the greater effect nuclear charge, the electrons in the $n=2$ shell in $\mathrm{F}$ or the $n=2$ shell in $\mathrm{B}$ ? Which will be closer to the nucleus?

Tracy Tourville
Tracy Tourville
Numerade Educator
01:34

Problem 18

Arrange the following atoms in order of increasing effective nuclear charge experienced by the electrons in the $n=2$ shell: Be, Br, Na, P, Se.

Tracy Tourville
Tracy Tourville
Numerade Educator
01:33

Problem 19

Which quantity must be determined experimentally in order to determine the bonding atomic radius of an atom? (a) The distance from the nucleus where the probability of finding an electron goes to zero. (b) The distance between the nuclei of two atoms that are bonded together. (c) The effective nuclear charge of an atom.

William Mills
William Mills
Numerade Educator
07:33

Problem 20

With the exception of helium, the noble gases condense to form solids when they are cooled sufficiently. At temperatures below $83 \mathrm{~K}$, argon forms a close-packed solid whose structure is shown below. (a) What is the apparent radius of an argon atom in solid argon, assuming the atoms touch as shown in this figure? (b) Is this value larger or smaller than the bonding atomic radius estimated for argon in Figure $7.7 ?$ (c) Based on this comparison would you say that the atoms are held together by chemical bonds in solid argon?

James Irizarry
James Irizarry
Numerade Educator
02:35

Problem 21

Tungsten has the highest melting point of any metal in the periodic table: $3422^{\circ} \mathrm{C}$. The distance between $\mathrm{W}$ atoms in tungsten metal is $274 \mathrm{pm}$. (a) What is the atomic radius of a tungsten atom in this environment? (This radius is called the metallic radius.) (b) If you put tungsten metal under high pressure, predict what would happen to the distance between $\mathrm{W}$ atoms.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
06:05

Problem 22

Which of the following statements about the bonding atomic radii in Figure 7.7 is incorrect? (i) For a given period, the radii of the representative elements generally decrease from left to right across a period. (ii) The radii of the representative elements for the $n=3$ period are all larger than those of the corresponding elements in the $n=2$ period. (iii) For most of the representative elements, the change in radius from the $n=2$ to the $n=3$ period is greater than the change in radius from $n=3$ to $n=4$. (iv) The radii of the transition elements generally increase moving from left to right within a period. (v) The large radii of the Group 1 elements are due to their relatively small effective nuclear charges.

Tracy Tourville
Tracy Tourville
Numerade Educator
03:11

Problem 23

Estimate the $\mathrm{P}-\mathrm{Cl}$ bond length from the data in Figure 7.7 and compare your value to the experimental $\mathrm{P}-\mathrm{Cl}$ bond length in phosphorus tetrachloride $\mathrm{PCl}_{3}, 204 \mathrm{pm}$.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:53

Problem 24

The experimental $\mathrm{Pb}-\mathrm{Cl}$ bond length in lead(II)chloride, $\mathrm{PbCl}_{2}$, is $244 \mathrm{pm}$. Based on this value and data in Figure 7.7 , predict the atomic radius of $\mathrm{Pb}$.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
04:11

Problem 25

Using only the periodic table, arrange each set of atoms in order from largest to smallest: $(\mathbf{a}) \mathrm{Ar}, \mathrm{As}, \mathrm{Kr} ;(\mathbf{b}) \mathrm{Cd}, \mathrm{Rb}, \mathrm{Te} ;(\mathbf{c})$ C, Cl, Cu.

Tracy Tourville
Tracy Tourville
Numerade Educator
02:55

Problem 26

Using only the periodic table, arrange each set of atoms in order of increasing radius: (a) Cs, Se, Te; (b) $\mathrm{S}, \mathrm{Si}, \mathrm{Sr} ;$ (c) P, Po, Pb.

Tracy Tourville
Tracy Tourville
Numerade Educator
01:47

Problem 27

Identify each statement as true or false: (a) Cations are larger than their corresponding neutral atoms. (b) $\mathrm{Li}^{+}$ is smaller than Li. (c) $\mathrm{Cl}^{-}$ is bigger than I $^{-}$.

William Mills
William Mills
Numerade Educator
02:55

Problem 28

Explain the following variations in atomic or ionic radii:
(a) $\mathrm{I}^{-}>\mathrm{I}>\mathrm{I}^{+}$
(b) $\mathrm{Ca}^{2+}>\mathrm{Mg}^{2+}>\mathrm{Be}^{2+}$
(c) $\mathrm{Fe}>\mathrm{Fe}^{2+}>\mathrm{Fe}^{3+}$

James Irizarry
James Irizarry
Numerade Educator
03:21

Problem 29

Which neutral atom is isoelectronic with each of the following ions? $\mathrm{H}^{-}, \mathrm{Ca}^{2+}, \mathrm{In}^{3+}, \mathrm{Ge}^{2+}$

Tracy Tourville
Tracy Tourville
Numerade Educator
01:58

Problem 30

Some ions do not have a corresponding neutral atom that has the same electron configuration. For each of the following ions, identify the neutral atom that has the same number of electrons and determine if this atom has the same electron configuration. (a) $\mathrm{Cl}^{-}$, (b) $\mathrm{Sc}^{3+}$, (c) $\mathrm{Fe}^{2+}$, (d) $\mathrm{Zn}^{2+}$, (e) $\mathrm{Sn}^{4+}$.

James Irizarry
James Irizarry
Numerade Educator
04:10

Problem 31

Consider the isoelectronic ions $\mathrm{F}^{-}$ and $\mathrm{Na}^{+}$. (a) Which ion is smaller? (b) Using Equation 7.1 and assuming that core electrons contribute 1.00 and valence electrons contribute 0.00 to the screening constant, $S$, calculate $Z_{\text {eff }}$ for the $2 p$ electrons in both ions. (c) Repeat this calculation using Slater's rules to estimate the screening constant, $S .(\mathbf{d})$ For isoelectronic ions, how are effective nuclear charge and ionic radius related?

William Mills
William Mills
Numerade Educator
13:09

Problem 32

Consider the isoelectronic ions $\mathrm{Cl}^{-}$ and $\mathrm{K}^{+}$. (a) Which ion is smaller? (b) Using Equation 7.1 and assuming that core electrons contribute 1.00 and valence electrons contribute nothing to the screening constant, $S,$ calculate $Z_{\text {eff }}$ for these two ions. (c) Repeat this calculation using Slater's rules to estimate the screening constant, $S .$ (d) For isoelectronic ions, how are effective nuclear charge and ionic radius related?

James Irizarry
James Irizarry
Numerade Educator
02:47

Problem 33

Consider $S, C l,$ and $K$ and their most common ions. (a) List the atoms in order of increasing size. (b) List the ions in order of increasing size. (c) Explain any differences in the orders of the atomic and ionic sizes.

William Mills
William Mills
Numerade Educator
05:31

Problem 34

Arrange each of the following sets of atoms and ions, in order of increasing size: (a) $\mathrm{Pb}, \mathrm{Pb}^{2+}, \mathrm{Pb}^{4+}$ (b) $\mathrm{V}^{3+}, \mathrm{Co}^{2+}, \mathrm{Co}^{3+}$ (c) $\mathrm{Se}^{2-}, \mathrm{S}^{2-}, \mathrm{Sn}^{2+} ;(\mathbf{d}) \mathrm{K}^{+}, \mathrm{Rb}^{+}, \mathrm{Br}^{-}$

Tracy Tourville
Tracy Tourville
Numerade Educator
05:28

Problem 35

Provide a brief explanation for each of the following: $(\mathbf{a}) \mathrm{Cl}^{-}$ is larger than Ar. (b) $\mathrm{P}^{3-}$ is larger than $\mathrm{S}^{2-}$. (c) $\mathrm{K}^{+}$ is larger than $\mathrm{Na}^{+} .(\mathbf{d}) \mathrm{F}^{-}$ is larger than $\mathrm{F}$.

Tracy Tourville
Tracy Tourville
Numerade Educator
13:24

Problem 36

In the ionic compounds LiF, $\mathrm{NaCl}, \mathrm{KBr},$ and RbI, the measured cation-anion distances are 201 pm (Li-F), 282 pm $(\mathrm{Na}-\mathrm{Cl}), 330 \mathrm{pm}(\mathrm{K}-\mathrm{Br}),$ and $367 \mathrm{pm}(\mathrm{Rb}-\mathrm{I}),$ respectively. $(\mathbf{a})$ Predict the cation-anion distance using the values of ionic radii given in Figure 7.8. (b) Calculate the difference between the experimentally measured ion-ion distances and the ones predicted from Figure $7.8 .$ (c) What estimates of the cationanion distance would you obtain for these four compounds using neutral atom bonding atomic radii? Are these estimates as accurate as the estimates using ionic radii?

James Irizarry
James Irizarry
Numerade Educator
03:08

Problem 37

Write equations that show the processes that describe the first, second, and third ionization energies of a chlorine atom. Which process would require the least amount of energy?

Tracy Tourville
Tracy Tourville
Numerade Educator
03:34

Problem 38

Write equations that show the process for (a) the first two ionization energies of zinc and (b) the fourth ionization energy of calcium.

Tracy Tourville
Tracy Tourville
Numerade Educator
01:52

Problem 39

Which element has the highest second ionization energy: Li, K, or Be?

William Mills
William Mills
Numerade Educator
05:36

Problem 40

Identify each statement as true or false: (a) Ionization energies are always endothermic. (b) Potassium has a larger first ionization energy than lithium. (c) The second ionization energy of the sodium atom is larger than the second ionization energy of the magnesium atom. (d) The third ionization energy is three times the first ionization energy of an atom.

Tracy Tourville
Tracy Tourville
Numerade Educator
02:12

Problem 41

(a) What is the general relationship between the size of an atom and its first ionization energy? (b) Which element in the periodic table has the largest ionization energy? Which has the smallest?

William Mills
William Mills
Numerade Educator
02:36

Problem 42

(a) What is the trend in first ionization energies as one proceeds down the group 17 elements? Explain how this trend relates to the variation in atomic radii. (b) What is the trend in first ionization energies as one moves across the fourth period from $\mathrm{K}$ to $\mathrm{Kr}$ ? How does this trend compare with the trend in atomic radii?

Tracy Tourville
Tracy Tourville
Numerade Educator
03:35

Problem 43

Based on their positions in the periodic table, predict which atom of the following pairs will have the smaller first ionization energy: (a) $\mathrm{Br}, \mathrm{Kr} ;$ (b) C, Ca; (c) $\mathrm{Li}, \mathrm{Rb} ;$ (d) $\mathrm{Pb}, \mathrm{Si} ;$ (e) $\mathrm{Al}, \mathrm{B}$.

Tracy Tourville
Tracy Tourville
Numerade Educator
03:08

Problem 44

For each of the following pairs, indicate which element has the smaller first ionization energy: (a) Cs, Cl; (b) Fe, Zn; (c) I, $\mathrm{Cl} ;(\mathbf{d}) \mathrm{Se}, \mathrm{Sn}$

Tracy Tourville
Tracy Tourville
Numerade Educator
05:04

Problem 45

Write the electron configurations for the following ions, and determine which have noble-gas configurations: (a) $\mathrm{Cu}^{2+}$, (b) $\mathrm{Ca}^{2+},(\mathbf{c}) \mathrm{N}^{3-}$, (d) $\mathrm{Ru}^{2+}$, (e) $\mathrm{H}^{-}$.

Tracy Tourville
Tracy Tourville
Numerade Educator
04:54

Problem 46

Write the electron configurations for the following ions, and determine which have noble-gas configurations: (a) $\mathrm{Ti}^{2+}$, (b) $\mathrm{Br}^{-}$, (c) $\mathrm{Mg}^{2+}$, (d) $\mathrm{Po}^{2-}$, (e) $\mathrm{Pt}^{2+},(\mathbf{f}) \mathrm{V}^{3+} .$

Tracy Tourville
Tracy Tourville
Numerade Educator
01:42

Problem 47

Give three examples of +2 ions that have an electron configuration of $n d^{10}(n=3,4,5 \ldots)$.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
01:49

Problem 48

Give examples of transition metal ions with +3 charge that have an electron configuration of $n d^{5}(n=3,4,5 \ldots)$.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:32

Problem 49

Write an equation for the first electron affinity of helium. Would you predict a positive or a negative energy value for this process? Is it possible to directly measure the first electron affinity of helium?

Tracy Tourville
Tracy Tourville
Numerade Educator
02:50

Problem 50

If the electron affinity for an element is a negative number, does it mean that the anion of the element is more stable than the neutral atom? Explain.

James Irizarry
James Irizarry
Numerade Educator
02:08

Problem 51

Which of the following, I or $\mathrm{I}^{-}$, will have a negative electron affinity?

Tracy Tourville
Tracy Tourville
Numerade Educator
01:46

Problem 52

What is the relationship between the ionization energy of an anion with a 1 - charge such as $\mathrm{F}^{-}$ and the electron affinity of the neutral atom, F?

James Irizarry
James Irizarry
Numerade Educator
03:59

Problem 53

Consider the first ionization energy of neon and the electron affinity of fluorine. (a) Write equations, including electron configurations, for each process. (b) These two quantities have opposite signs. Which will be positive, and which will be negative? (c) Would you expect the magnitudes of these two quantities to be equal? If not, which one would you expect to be larger?

William Mills
William Mills
Numerade Educator
04:46

Problem 54

Consider the following equation:
$$
\mathrm{Al}^{3+}(g)+e^{-} \longrightarrow \mathrm{Al}^{2+}(g)
$$
Which of the following statements are true? (i) The energy change for this process is the second electron affinity of $\mathrm{Al}$ atom since $\mathrm{Al}^{2+}(g)$ is formed. (ii) The energy change for this process is the negative of the third ionization energy of the Al atom. (iii) The energy change for this process is the electron affinity of the $\mathrm{Al}^{2+}$ ion.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:36

Problem 55

(a) Does metallic character increase, decrease, or remain unchanged as one goes from left to right across a row of the periodic table? (b) Does metallic character increase, decrease, or remain unchanged as one goes down a column of the periodic table? (c) Are the periodic trends in (a) and (b) the same as or different from those for first ionization energy?

William Mills
William Mills
Numerade Educator
01:11

Problem 56

You read the following statement about two elements $X$ and Y: One of the elements is a good conductor of electricity, and the other is a semiconductor. Experiments show that the first ionization energy of $X$ is twice as great as that of $Y .$ Which element has the greater metallic character?

James Irizarry
James Irizarry
Numerade Educator
00:52

Problem 57

Discussing this chapter, a classmate says, "An element that commonly forms a cation is a metal." Do you agree or disagree?

William Mills
William Mills
Numerade Educator
02:50

Problem 58

Discussing this chapter, a classmate says, "Since elements that form cations are metals and elements that form anions are nonmetals, elements that do not form ions are metalloids." Do you agree or disagree?

James Irizarry
James Irizarry
Numerade Educator
02:12

Problem 59

Predict whether each of the following oxides is ionic or molecular: $\mathrm{ZnO}, \mathrm{K}_{2} \mathrm{O}, \mathrm{SO}_{2}, \mathrm{OF}_{2}, \mathrm{TiO}_{2}$

Tracy Tourville
Tracy Tourville
Numerade Educator
01:56

Problem 60

Some metal oxides, such as $\mathrm{Sc}_{2} \mathrm{O}_{3},$ do not react with pure water, but they do react when the solution becomes either acidic or basic. Do you expect $\mathrm{Sc}_{2} \mathrm{O}_{3}$ to react when the solution becomes acidic or when it becomes basic? Write a balanced chemical equation to support your answer.

James Irizarry
James Irizarry
Numerade Educator
02:04

Problem 61

Would you expect zirconium(II) oxide, $\mathrm{ZrO},$ to react more readily with $\mathrm{HCl}(a q)$ or $\mathrm{NaOH}(a q) ?$

Tracy Tourville
Tracy Tourville
Numerade Educator
02:19

Problem 62

Arrange the following oxides in order of increasing acidity: $\mathrm{K}_{2} \mathrm{O}, \mathrm{BaO}, \mathrm{ZnO}, \mathrm{H}_{2} \mathrm{O}, \mathrm{CO}_{2}, \mathrm{SO}_{2}$.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
03:10

Problem 63

Chlorine reacts with oxygen to form $\mathrm{Cl}_{2} \mathrm{O}_{7} .(\mathbf{a})$ What is the name of this product (see Table 2.6$) ?(\mathbf{b})$ Write a balanced equation for the formation of $\mathrm{Cl}_{2} \mathrm{O}_{7}(l)$ from the elements. (c) Would you expect $\mathrm{Cl}_{2} \mathrm{O}_{7}$ to be more reactive toward $\mathrm{H}^{+}(a q)$ or $\mathrm{OH}^{-}(a q) ?(\mathbf{d})$ If the oxygen in $\mathrm{Cl}_{2} \mathrm{O}_{7}$ is considered to have the -2 oxidation state, what is the oxidation state of the Cl? What is the electron configuration of Cl in this oxidation state?

Kevin Chimex
Kevin Chimex
Numerade Educator
04:55

Problem 64

An element X reacts with oxygen to form $\mathrm{XO}_{2}$ and with chlorine to form $\mathrm{XCl}_{4} . \mathrm{XO}_{2}$ is a white solid that melts at high temperatures (above $1000^{\circ} \mathrm{C}$ ). Under usual conditions, $\mathrm{XCl}_{4}$ is a colorless liquid with a boiling point of $58^{\circ} \mathrm{C}$. (a) $\mathrm{XCl}_{4}$ reacts with water to form $\mathrm{XO}_{2}$ and another product. What is the likely identity of the other product? (b) Do you think that element $X$ is a metal, nonmetal, or metalloid? (c) By using a sourcebook such as the CRC Handbook of Chemistry and Physics, try to determine the identity of element $\mathrm{X}$.

James Irizarry
James Irizarry
Numerade Educator
04:41

Problem 65

Write balanced equations for the following reactions: (a) boron trichloride with water, $(\mathbf{b})$ cobalt (II) oxide with nitric acid, (c) phosphorus pentoxide with water, (d) carbon dioxide with aqueous barium hydroxide.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
03:09

Problem 66

Write balanced equations for the following reactions: (a) sulfur dioxide with water, (b) lithium oxide in water, $(\mathbf{c})$ zinc oxide with dilute hydrochloric acid, $(\mathbf{d})$ arsenic trioxide with aqueous potassium hydroxide.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:22

Problem 67

(a) Why is calcium generally more reactive than beryllium? (b) Why is calcium generally less reactive than rubidium?

Tracy Tourville
Tracy Tourville
Numerade Educator
04:34

Problem 68

Copper and calcium both form +2 ions, but copper is far less reactive. Suggest an explanation, taking into account the ground-state electron configurations of these elements and their atomic radii.

Tracy Tourville
Tracy Tourville
Numerade Educator
02:08

Problem 69

Write a balanced equation for the reaction that occurs in each of the following cases: (a) Potassium metal is exposed to an atmosphere of chlorine gas. (b) Strontium oxide is added to water. (c) A fresh surface of lithium metal is exposed to oxygen gas. (d) Sodium metal reacts with molten sulfur.

Kevin Chimex
Kevin Chimex
Numerade Educator
02:41

Problem 70

Write a balanced equation for the reaction that occurs in each of the following cases: (a) Lithium is added to water. (b) Calcium is added to water. (c) Potassium reacts with chlorine gas. (d) Rubidium reacts with oxygen.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
03:24

Problem 71

(a) As described in Section 7.7, the alkali metals react with hydrogen to form hydrides and react with halogens to form halides. Compare the roles of hydrogen and halogens in these reactions. Write balanced equations for the reaction of fluorine with calcium and for the reaction of hydrogen with calcium. (b) What is the oxidation number and electron configuration of calcium in each product?

Kevin Chimex
Kevin Chimex
Numerade Educator
08:20

Problem 72

Potassium and hydrogen react to form the ionic compound potassium hydride. (a) Write a balanced equation for this reaction. (b) Use data in Figures 7.10 and 7.12 to determine the energy change in $\mathrm{kJ} / \mathrm{mol}$ for the following two reactions:
$$
\begin{array}{l}
\mathrm{K}(g)+\mathrm{H}(g) \longrightarrow \mathrm{K}^{+}(g)+\mathrm{H}^{-}(g) \\
\mathrm{K}(g)+\mathrm{H}(g) \longrightarrow \mathrm{K}^{-}(g)+\mathrm{H}^{+}(g)
\end{array}
$$
(c) Based on your calculated energy changes in (b), which of these reactions is energetically more favorable (or less unfavorable)? (d) Is your answer to (c) consistent with the description of potassium hydride as containing hydride ions?

James Irizarry
James Irizarry
Numerade Educator
03:39

Problem 73

Compare the elements bromine and chlorine with respect to the following properties: (a) electron configuration, (b) most common ionic charge, $(\mathbf{c})$ first ionization energy, $(\mathbf{d})$ reactivity toward water, $(\mathbf{e})$ electron affinity, $(\mathbf{f})$ atomic radius. Account for the differences between the two elements.

Kevin Chimex
Kevin Chimex
Numerade Educator
03:54

Problem 74

Little is known about the properties of astatine, At, because of its rarity and high radioactivity. Nevertheless, it is possible for us to make many predictions about its properties. (a) Do you expect the element to be a gas, liquid, or solid at room temperature? Explain. (b) Would you expect At to be a metal, nonmetal, or metalloid? Explain. (c) What is the chemical formula of the compound it forms with Na?

James Irizarry
James Irizarry
Numerade Educator
01:58

Problem 75

Until the early 1960 s, the group 18 elements were called the inert gases. (a) Why was the term inert gases dropped? (b) What discovery triggered this change in name? (c) What name is applied to the group now?

Tracy Tourville
Tracy Tourville
Numerade Educator
07:53

Problem 76

(a) Why does xenon react with fluorine, whereas neon does not? (b) Using appropriate reference sources, look up the bond lengths of Xe-F bonds in several molecules. How do these numbers compare to the bond lengths calculated from the atomic radii of the elements?

James Irizarry
James Irizarry
Numerade Educator
02:52

Problem 77

Write a balanced equation for the reaction that occurs in each of the following cases: (a) White phorphrous, $\mathrm{P}_{4}(\mathrm{~s})$ reacts with chlorine gas. (b) Sodium metal reacts with water. (c) Hydrogen bromide gas reacts with chlorine gas. (d) Aluminum trichloride reacts with aqueous sodium hydroxide.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
03:52

Problem 78

Write a balanced equation for the reaction that occurs in each of the following cases: (a) Calcium metal is heated in an atmosphere of oxygen gas. (b) Copper oxide is heated in an atmosphere of hydrogen gas. (c) Chlorine reacts with nitrogen gas. (d) Boron tribromide reacts with water.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
03:07

Problem 79

Consider the stable elements through lead $(Z=82) .$ In how many instances are the atomic weights of the elements out of order relative to the atomic numbers of the elements?

Kevin Chimex
Kevin Chimex
Numerade Educator
05:01

Problem 80

Figure 7.4 shows the radial probability distribution functions for the $2 s$ orbitals and $2 p$ orbitals. (a) Which orbital, $2 s$ or $2 p,$ has more electron density close to the nucleus? (b) How would you modify Slater's rules to adjust for the difference in electronic penetration of the nucleus for the $2 s$ and $2 p$ orbitals?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:26

Problem 81

(a) If the core electrons were totally effective at screening the valence electrons and the valence electrons provided no screening for each other, what would be the effective nuclear charge acting on the $3 s$ and $3 p$ valence electrons in P? (b) Repeat these calculations using Slater's rules. $(\mathbf{c})$ Detailed calculations indicate that the effective nuclear charge is $5.6+$ for the $3 s$ electrons and $4.9+$ for the $3 p$ electrons. Why are the values for the $3 s$ and $3 p$ electrons different? (d) If you remove a single electron from a P atom, which orbital will it come from?

Lottie Adams
Lottie Adams
Numerade Educator
05:26

Problem 82

As we move across a period of the periodic table, why do the sizes of the transition elements change more gradually than those of the representative elements?

James Irizarry
James Irizarry
Numerade Educator
07:00

Problem 83

In the series of group 16 hydrides, of general formula $\mathrm{MH}_{2}$, the measured bond distances are $\mathrm{O}-\mathrm{H}, 96 \mathrm{pm} ;$ $\mathrm{S}-\mathrm{H}, 134 \mathrm{pm} ;$ Se? H, $146 \mathrm{pm} .$ (a) Compare these values with those estimated by use of the atomic radii in Figure 7.7. (b) Explain the steady increase in $\mathrm{M}-\mathrm{H}$ bond distance in this series in terms of the electron configurations of the $\mathrm{M}$ atoms.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
01:35

Problem 84

In Table 7.8 , the bonding atomic radius of neon is listed as $58 \mathrm{pm}$, whereas that for xenon is listed as $140 \mathrm{pm}$. A classmate of yours states that the value for Xe is more realistic than the one for Ne. Is she correct? If so, what is the basis for her statement?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:20

Problem 85

The $\mathrm{P}-\mathrm{P}$ bond length in white phosphorus is $189 \mathrm{pm}$. The $\mathrm{Cl}-\mathrm{Cl}$ bond length in $\mathrm{Cl}_{2}$ is $199 \mathrm{pm}$. (a) Based on these data, what is the predicted $\mathrm{P}-\mathrm{Cl}$ bond length in phosphorus trichloride, $\mathrm{PCl}_{3}$, in which each of the three $\mathrm{Cl}$ atoms is bonded to the P atom? (b) What bond length is predicted for $\mathrm{PCl}_{3}$, using the atomic radii in Figure 7.7 ?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:10

Problem 86

The following observations are made about two hypothetical elements $A$ and $B$ : The $A-A$ and $B-B$ bond lengths in the elemental forms of A and B are 236 and $194 \mathrm{pm}$, respectively. A and B react to form the binary compound $\mathrm{AB}_{2}$, which has a linear structure (that is $\left.\angle \mathrm{B}-\mathrm{A}-\mathrm{B}=180^{\circ}\right)$. Based on these statements, predict the separation between the two B nuclei in a molecule of $\mathrm{AB}_{2}$.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
03:13

Problem 87

Elements in group 17 in the periodic table are called the halogens; elements in group 16 are called the chalcogens. (a) What is the most common oxidation state of the chalcogens compared to the halogens? (b) For each of the following periodic properties, state whether the halogens or the chalcogens have larger values: atomic radii, ionic radii of the most common oxidation state, first ionization energy, second ionization energy.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
03:54

Problem 88

Note from the following table that there is a significant increase in atomic radius upon moving from $\mathrm{Y}$ to La, whereas the radii of Zr to Hf are the same. Suggest an explanation for this effect.
$$
\begin{aligned}
&\text { Atomic Radii (pm) }\\
&\begin{array}{cccc}
\hline \text { Sc } & 170 & \text { Ti } & 160 \\
\text { Y } & 190 & \text { Zr } & 175 \\
\text { La } & 207 & \text { Hf } & 175 \\
\hline
\end{array}
\end{aligned}
$$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
08:33

Problem 89

(a) Which ion is smaller, $\mathrm{Co}^{3+}$ or $\mathrm{Co}^{4+} ?(\mathbf{b})$ In a lithium-ion battery that is discharging to power a device, for every $\mathrm{Li}^{+}$ that inserts into the lithium cobalt oxide electrode, a $\mathrm{Co}^{4+}$ ion must be reduced to a $\mathrm{Co}^{3+}$ ion to balance charge. Using the CRC Handbook of Chemistry and Physics or other standard reference, find the ionic radii of $\mathrm{Li}^{+}, \mathrm{Co}^{3+},$ and $\mathrm{Co}^{4+} .$ Order these ions from smallest to largest. (c) Will the lithium cobalt oxide cathode expand or contract as lithium ions are inserted? (d) Lithium is not nearly as abundant as sodium. If sodium ion batteries were developed that function in the same manner as lithium ion batteries, do you think "sodium cobalt oxide" would still work as the electrode material? Explain. (e) If you don't think cobalt would work as the redox-active partner ion in the sodium version of the electrode, suggest an alternative metal ion and explain your reasoning.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
09:03

Problem 90

The ionic substance strontium oxide, $\mathrm{SrO},$ forms from the reaction of strontium metal with molecular oxygen. The arrangement of the ions in solid $\mathrm{SrO}$ is analogous to that in solid $\mathrm{NaCl}:$ (a) Write a balanced equation for the formation of $\operatorname{SrO}(s)$ from its elements. (b) Based on the ionic radii in Figure 7.8 , predict the length of the side of the cube in the figure (the distance from the center of an atom at one corner to the center of an atom at a neighboring corner). (c) The density of $\mathrm{SrO}$ is $5.10 \mathrm{~g} / \mathrm{cm}^{3}$. Given your answer to part (b), how many formula units of $\mathrm{SrO}$ are contained in the cube shown here?

James Irizarry
James Irizarry
Numerade Educator
06:17

Problem 91

Explain the variation in the ionization energies of carbon, as displayed in this graph:

Kevin Chimex
Kevin Chimex
Numerade Educator
08:30

Problem 92

Group 14 elements have much more negative electron affinities than their neighbors in groups 13 and 15 (see Figure 7.12). Which of the following statements best explains this observation? (i) The group 14 elements have much higher first ionization energies than their neighbors in groups 13 and 15. (ii) The addition of an electron to a group 14 element leads to a half-filled $n p^{3}$ outer electron configuration. (iii) The group 14 elements have unusually large atomic radii. (iv) The group 14 elements are easier to vaporize than are the group 13 and 15 elements.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
04:39

Problem 93

In the chemical process called electron transfer, an electron is transferred from one atom or molecule to another. (We will talk about electron transfer extensively in Chapter
20.) A simple electron transfer reaction is
$$
\mathrm{A}(g)+\mathrm{A}(g) \longrightarrow \mathrm{A}^{+}(g)+\mathrm{A}^{-}(g)
$$
In terms of the ionization energy and electron affinity of atom A, what is the energy change for this reaction? For a representative nonmetal such as chlorine, is this process exothermic? For a representative metal such as sodium, is this process exothermic?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
03:48

Problem 94

(a) Use orbital diagrams to illustrate what happens when an Oxygen atom gains two electrons. (b) Why does $\mathrm{O}^{3-}$ not exist?

James Irizarry
James Irizarry
Numerade Educator
12:27

Problem 95

Use electron configurations to explain the following observations: $(\mathbf{a})$ The first ionization energy of phosphorus is greater than that of sulfur. (b) The electron affinity of nitrogen is lower (less negative) than those of both carbon and oxygen. (c) The second ionization energy of oxygen is greater than the first ionization energy of fluorine. (d) The third ionization energy of manganese is greater than those of both chromium and iron.

Dr. Satish Ingale
Dr. Satish Ingale
Numerade Educator
03:38

Problem 96

Identify a +2 cation that has the following ground state electron configurations: (a) $[\mathrm{Ne}]$ (b) $[\mathrm{Ar}] 3 d^{9}$ (c) $[\mathrm{Xe}] 4 \mathrm{f}^{14} 5 d^{10} 6 s^{2}$.

Tracy Tourville
Tracy Tourville
Numerade Educator
02:42

Problem 97

Which of the following chemical equations is connected to the definitions of (a) the first ionization energy of oxygen,
(b) the second ionization energy of ox ' ygen, and $(\mathbf{c})$ the electron affinity of oxygen?
(i) $\mathrm{O}(g)+\mathrm{e}^{-} \longrightarrow \mathrm{O}^{-}(g)$
(ii) $\mathrm{O}(g) \longrightarrow \mathrm{O}^{+}(g)+\mathrm{e}^{-}$
(iii) $\mathrm{O}(g)+2 \mathrm{e}^{-} \longrightarrow \mathrm{O}^{2-}(g)$
(iv) $\mathrm{O}(g) \longrightarrow \mathrm{O}^{2+}(g)+2 \mathrm{e}^{-}$ $(\mathbf{v}) \mathrm{O}^{+}(g) \longrightarrow \mathrm{O}^{2+}(g)+\mathrm{e}^{-}$

Tracy Tourville
Tracy Tourville
Numerade Educator
01:43

Problem 98

The electron affinities, in $\mathrm{kJ} / \mathrm{mol}$, for the group 11 and group 12 metals are as follows:
$$
\begin{array}{|c|l|}
\hline \mathrm{Cu} & \mathrm{Zn} \\
-119 & >0 \\
\hline \mathrm{Ag} & \mathrm{Cd} \\
-126 & >0 \\
\hline \mathrm{Au} & \mathrm{Hg} \\
-223 & >0 \\
\hline
\end{array}
$$
(a) Why are the electron affinities of the group 12 elements greater than zero? (b) Why do the electron affinities of the group 11 elements become more negative as we move down the group?

Lottie Adams
Lottie Adams
Numerade Educator
08:56

Problem 99

Hydrogen is an unusual element because it behaves in some ways like the alkali metal elements and in other ways like nonmetals. Its properties can be explained in part by its electron configuration and by the values for its ionization energy and electron affinity. (a) Explain why the electron affinity of hydrogen is much closer to the values for the alkali elements than for the halogens. (b) Is the following statement true? "Hydrogen has the smallest bonding atomic radius of any element that forms chemical compounds." If not, correct it. If it is, explain in terms of electron configurations. (c) Explain why the ionization energy of hydrogen is closer to the values for the halogens than for the alkali metals. (d) The hydride ion is $\mathrm{H}^{-}$. Write out the process corresponding to the first ionization energy of the hydride ion. (e) How does the process in part (d) compare to the process for the electron affinity of a neutral hydrogen atom?

Kevin Chimex
Kevin Chimex
Numerade Educator
02:22

Problem 100

The first ionization energy of the oxygen molecule is the energy required for the following process:
$$
\mathrm{O}_{2}(g) \longrightarrow \mathrm{O}_{2}^{+}(g)+\mathrm{e}^{-}
$$
The energy needed for this process is $1175 \mathrm{~kJ} / \mathrm{mol}$, very similar to the first ionization energy of Xe. Would you expect $\mathrm{O}_{2}$ to react with $\mathrm{F}_{2} ?$ If so, suggest a product or products of this reaction.

James Irizarry
James Irizarry
Numerade Educator
06:44

Problem 101

It is possible to define metallic character as we do in this book and base it on the reactivity of the element and the ease with which it loses electrons. Alternatively, one could measure how well electricity is conducted by each of the elements to determine how "metallic" the elements are. On the basis of conductivity, there is not much of a trend in the periodic table: Silver is the most conductive metal, and manganese the least. Look up the first ionization energies of silver and manganese; which of these two elements would you call more metallic based on the way we define it in this book?

Dr. Satish Ingale
Dr. Satish Ingale
Numerade Educator
02:04

Problem 102

Which of the following is the expected product of the reaction of $\mathrm{Mg}(s)$ and $\mathrm{N}_{2}(g)$ under heat? (i) $\mathrm{Mg}_{3} \mathrm{~N}(s),(\mathbf{i i}) \mathrm{MgN}_{2}(s),$ (iii) $\mathrm{Mg}_{3} \mathrm{~N}_{2}(s)$ (iv) $\mathrm{Mg}(s)$ and $\mathrm{N}_{2}(g)$ will not react with one another.

Tracy Tourville
Tracy Tourville
Numerade Educator
03:39

Problem 103

Elemental barium reacts more violently with water than does elemental calcium. Which of the following best explains this difference in reactivity? (i) Calcium has greater metallic character than does barium. (ii) The electron affinity of calcium is smaller than that of barium. (iii) The first and second ionization energies of barium are less than those of calcium. (iv) The atomic radius of barium is smaller than that of calcium. ( $\mathbf{v}$ ) The ionic radius of the barium ion is larger than that of the calcium ion.

Tracy Tourville
Tracy Tourville
Numerade Educator
06:30

Problem 104

(a) One of the alkali metals reacts with oxygen to form a solid white substance. When this substance is dissolved in water, the solution gives a positive test for hydrogen peroxide, $\mathrm{H}_{2} \mathrm{O}_{2}$. When the solution is tested in a burner flame, a lilac-purple flame is produced. What is the likely identity of the metal? (b) Write a balanced chemical equation for the reaction of the white substance with water.

James Irizarry
James Irizarry
Numerade Educator
09:33

Problem 105

Zinc in its $2+$ oxidation state is an essential metal ion for life. $\mathrm{Zn}^{2+}$ is found bound to many proteins that are involved in biological processes, but unfortunately $\mathrm{Zn}^{2+}$ is hard to detect by common chemical methods. Therefore, scientists who are interested in studying $\mathrm{Zn}^{2+}$ -containing proteins frequently substitute $\mathrm{Cd}^{2+}$ for $\mathrm{Zn}^{2+},$ since $\mathrm{Cd}^{2+}$ is easier to detect. $(\mathbf{a}) \mathrm{On}$ the basis of the properties of the elements and ions discussed in this chapter and their positions in the periodic table, describe the pros and cons of using $\mathrm{Cd}^{2+}$ as a $\mathrm{Zn}^{2+}$ substitute. (b) Proteins that speed up (catalyze) chemical reactions are called enzymes. Many enzymes are required for proper metabolic reactions in the body. One problem with using $\mathrm{Cd}^{2+}$ to replace $\mathrm{Zn}^{2+}$ in enzymes is that $\mathrm{Cd}^{2+}$ substitution can decrease or even eliminate enzymatic activity. Can you suggest a different metal ion that might replace $\mathrm{Zn}^{2+}$ in enzymes instead of $\mathrm{Cd}^{2+}$ ? Justify your answer.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
00:57

Problem 106

A historian discovers a nineteenth-century notebook in which some observations, dated $1822,$ were recorded on a substance thought to be a new element. Here are some of the data recorded in the notebook: "Ductile, silver-white, metallic looking. Softer than lead. Unaffected by water. Stable in air. Melting point: $153^{\circ} \mathrm{C}$. Density: $7.3 \mathrm{~g} / \mathrm{cm}^{3}$. Electrical conductivity: $20 \%$ that of copper. Hardness: About $1 \%$ as hard as iron. When $4.20 \mathrm{~g}$ of the unknown is heated in an excess of oxygen, $5.08 \mathrm{~g}$ of a white solid is formed. The solid could be sublimed by heating to over $800^{\circ} \mathrm{C} . "($ a) Using information in the text and the CRC Handbook of Chemistry and Physics, and making allowances for possible variations in numbers from current values, identify the element reported. (b) Write a balanced chemical equation for the reaction with oxygen. (c) Judging from Figure 7.1, might this nineteenth-century investigator have been the first to discover a new element?

Lottie Adams
Lottie Adams
Numerade Educator
05:15

Problem 107

In April 2010 , a research team reported that it had made Element 117 . This discovery was confirmed in 2012 by additional experiments. Write the ground-state electron configuration for Element 117 and estimate values for its first ionization energy, electron affinity, atomic size, and common oxidation state based on its position in the periodic table.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
07:06

Problem 108

We will see in Chapter 12 that semiconductors are materials that conduct electricity better than nonmetals but not as well as metals. The only two elements in the periodic table that are technologically useful semiconductors are silicon and germanium. Integrated circuits in computer chips today are based on silicon. Compound semiconductors are also used in the electronics industry. Examples are gallium arsenide, GaAs; gallium phosphide, GaP; cadmium sulfide, CdS; and cadmium selenide, CdSe. (a) What is the relationship between the compound semiconductors' compositions and the positions of their elements on the periodic table relative to Si and Ge? (b) Workers in the semiconductor industry refer to "II-VI" and "III-V" materials, using Roman numerals. Can you identify which compound semiconductors are II-VI and which are III-V? (c) Suggest other compositions of compound semiconductors based on the positions of their elements in the periodic table.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
13:44

Problem 109

Moseley established the concept of atomic number by studying $X$ rays emitted by the elements. The $X$ rays emitted by some of the elements have the following wavelengths:
$$
\begin{array}{cc}
\hline \text { Element } & \text { Wavelength (pm) } \\
\hline \text { Ne } & 1461 \\
\text { Ca } & 335.8 \\
\text { Zn } & 143.5 \\
\text { Zr } & 78.6 \\
\text { Sn } & 49.1 \\
\hline
\end{array}
$$
(a) Calculate the frequency, $\nu,$ of the $X$ rays emitted by each of the elements, in Hz. (b) Plot the square root of $\nu$ versus the atomic number of the element. What do you observe about the plot? (c) Explain how the plot in part (b) allowed Moseley to predict the existence of undiscovered elements. (d) Use the result from part (b) to predict the X-ray wavelength emitted by iron. (e) A particular element emits X rays with a wavelength of $98.0 \mathrm{pm}$. What element do you think it is?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
11:42

Problem 110

(a) Write the electron configuration for $\mathrm{Li}$ and estimate the effective nuclear charge experienced by the valence electron. (b) The energy of an electron in a one-electron atom or ion equals $\left(-2.18 \times 10^{-18} \mathrm{~J}\right)\left(\frac{Z^{2}}{n^{2}}\right),$ where $Z$ is the nuclear charge and $n$ is the principal quantum number of the electron. Estimate the first ionization energy of Li. (c) Compare the result of your calculation with the value reported in Table 7.4 and explain the difference. (d) What value of the effective nuclear charge gives the proper value for the ionization energy? Does this agree with your explanation in part (c)?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
08:50

Problem 111

One way to measure ionization energies is ultraviolet photoelectron spectroscopy (PES), a technique based on the photoelectric effect. eoo (Section 6.2) In PES, monochromatic light is directed onto a sample, causing electrons to be emitted. The kinetic energy of the emitted electrons is measured. The difference between the energy of the photons and the kinetic energy of the electrons corresponds to the energy needed to remove the electrons (that is, the ionization energy). Suppose that a PES experiment is performed in which mercury vapor is irradiated with ultraviolet light of wavelength $58.4 \mathrm{nm} .(\mathbf{a})$ What is the energy of a photon of this light, in joules? (b) Write an equation that shows the process corresponding to the first ionization energy of $\mathrm{Hg}$. (c) The kinetic energy of the emitted electrons is measured to be $1.72 \times 10^{-18} \mathrm{~J}$. What is the first ionization energy of $\mathrm{Hg}$, in $\mathrm{kJ} / \mathrm{mol} ?$ (d) Using Figure 7.10 , determine which of the halogen elements has a first ionization energy closest to that of mercury.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
21:59

Problem 112

Mercury in the environment can exist in oxidation states 0 , + 1 , and +2 . One major question in environmental chemistry research is how to best measure the oxidation state of mercury in natural systems; this is made more complicated by the fact that mercury can be reduced or oxidized on surfaces differently than it would be if it were free in solution. XPS, X-ray photoelectron spectroscopy, is a technique related to PES (see Exercise 7.111 ), but instead of using ultraviolet light to eject valence electrons, X rays are used to eject core electrons. The energies of the core electrons are different for different oxidation states of the element. In one set of experiments, researchers examined mercury contamination of minerals in water. They measured the XPS signals that corresponded to electrons ejected from mercury's $4 f$ orbitals at $105 \mathrm{eV},$ from an $\mathrm{X}$ -ray source that provided $1253.6 \mathrm{eV}$ of energy $\left(1 \mathrm{ev}=1.602 \times 10^{-19} \mathrm{~J}\right)$ The oxygen on the mineral surface gave emitted electron energies at $531 \mathrm{eV}$, corresponding to the $1 s$ orbital of oxygen. Overall the researchers concluded that oxidation states were +2 for $\mathrm{Hg}$ and -2 for O. (a) Calculate the wavelength of the $X$ rays used in this experiment. (b) Compare the energies of the $4 f$ electrons in mercury and the $1 s$ electrons in oxygen from these data to the first ionization energies of mercury and oxygen from the data in this chapter. (c) Write out the ground-state electron configurations for $\mathrm{Hg}^{2+}$ and $\mathrm{O}^{2-}$; which electrons are the valence electrons in each case?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
20:21

Problem 113

When magnesium metal is burned in air (Figure 3.6), two products are produced. One is magnesium oxide, $\mathrm{MgO}$. The other is the product of the reaction of $\mathrm{Mg}$ with molecular nitrogen, magnesium nitride. When water is added to magnesium nitride, it reacts to form magnesium oxide and ammonia gas. (a) Based on the charge of the nitride ion (Table 2.5), predict the formula of magnesium nitride. (b) Write a balanced equation for the reaction of magnesium nitride with water. What is the driving force for this reaction? (c) In an experiment, a piece of magnesium ribbon is burned in air in a crucible. The mass of the mixture of $\mathrm{MgO}$ and magnesium nitride after burning is $0.470 \mathrm{~g}$. Water is added to the crucible, further reaction occurs, and the crucible is heated to dryness until the final product is $0.486 \mathrm{~g}$ of $\mathrm{MgO}$. What was the mass percentage of magnesium nitride in the mixture obtained after the initial burning? (d) Magnesium nitride can also be formed by reaction of the metal with ammonia at high temperature. Write a balanced equation for this reaction. If a 6.3-g Mg ribbon reacts with $2.57 \mathrm{~g} \mathrm{NH}_{3}(g)$ and the reaction goes to completion, which component is the limiting reactant? What mass of $\mathrm{H}_{2}(g)$ is formed in the reaction? (e) The standard enthalpy of formation of solid magnesium nitride is $-461.08 \mathrm{~kJ} / \mathrm{mol}$. Calculate the standard enthalpy change for the reaction between magnesium metal and ammonia gas.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
10:44

Problem 114

(a) The measured $\mathrm{Bi}-\mathrm{Br}$ bond length in bismuth tribromide, $\mathrm{BiBr}_{3},$ is $263 \mathrm{pm} .$ Based on this value and the data in Figure 7.8 predict the atomic radius of Bi. (b) Bismuth tribromide is soluble in acidic solution. It is formed by treating solid bismuth(III) oxide with aqueous hydrobromic acid. Write a balanced chemical equation for this reaction. (c) While bismuth(III) oxide is soluble in acidic solutions, it is insoluble in basic solutions such as $\mathrm{NaOH}(a q)$. Based on these properties, is bismuth characterized as a metallic, metalloid, or nonmetallic element? (d) Treating bismuth with fluorine gas forms $\mathrm{BiF}_{5}$. Use the electron configuration of Bi to explain the formation of a compound with this formulation. (e) While it is possible to form $\mathrm{BiF}_{5}$ in the manner just described, pentahalides of bismuth are not known for the other halogens. Explain why the pentahalide might form with fluorine but not with the other halogens. How does the behavior of bismuth relate to the fact that xenon reacts with fluorine to form compounds but not with the other halogens?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:51

Problem 115

Potassium superoxide, $\mathrm{KO}_{2},$ is often used in oxygen masks (such as those used by firefighters) because $\mathrm{KO}_{2}$ reacts with $\mathrm{CO}_{2}$ to release molecular oxygen. Experiments indicate that 2 mol of $\mathrm{KO}_{2}(s)$ react with each mole of $\mathrm{CO}_{2}(g) .(\mathbf{a})$ The products of the reaction are $\mathrm{K}_{2} \mathrm{CO}_{3}(s)$ and $\mathrm{O}_{2}(g) .$ Write a balanced equation for the reaction between $\mathrm{KO}_{2}(s)$ and $\mathrm{CO}_{2}(g) .(\mathbf{b})$ Indicate the oxidation number for each atom involved in the reaction in part (a). What elements are being oxidized and reduced? (c) What mass of $\mathrm{KO}_{2}(s)$ is needed to consume $18.0 \mathrm{~g} \mathrm{CO}_{2}(g)$ ? What mass of $\mathrm{O}_{2}(g)$ is produced during this reaction?

Kevin Chimex
Kevin Chimex
Numerade Educator