Kelter P., Mosher M., Scott A. Chemistry. The Practical Science

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FIGURE 7.13

Electronegativity values for selected ele-

ments. Electronegativity increases as we

go up and to the right in the periodic

table.

However, our bodies require different amounts of each vitamin as part of our

daily diet. Why? Vitamin A is a fat-soluble compound that our bodies retain.

Vitamin C, on the other hand, is a water-soluble compound that is readily ex-

creted from our bodies. This means that we must consume vitamin C in larger

quantities than vitamin A and on a more regular basis. As we will discuss in the

next chapter, electronegativities are the basis for understanding the solubility of

these vitamins.

Pauling derived the electronegativity values by comparing the energies re-

quired to break various bonds, as we will discuss more fully in Chapter 8. The

values shown in Figure 7.13 reveal two general trends. The ﬁrst of these trends is

that electronegativities generally increase as we move along periods all the way to

Group VIIA. The second trend indicates that electronegativities generally decrease

as we move down groups. This means that the lowest electronegativity values are

at the bottom of Group IA and that the element ﬂuorine has the highest.

Do these

trends make sense?

The increase in electronegativity along periods occurs be-

cause the atoms are progressively smaller. In addition, atoms on the right-hand

288 Chapter 7 Periodic Properties of the Elements

1.1

2.1

1.0

1.5

0.9

1.2

0.8

1.0

0.8

1.0

0.7

0.9

0.7

0.9

1.3

1.2

1.1

1.5

1.4

1.3

1.6

1.5

1.6

1.8

1.7

1.5

1.9

1.8

2.2

1.9

2.2

1.9

2.2

1.9

2.4

1.6

1.7

1.9

1.6

1.7

1.8

1.5

2.0

1.8

1.9

1.8

2.5

2.0

1.9

2.1

3.0

2.4

2.1

2.0

2.5

3.5

2.8

2.5

2.2

3.0

4.0

Increasing electronegativity

Decreasing electronegativity

1.3

1.2

1.1

1.3

1.4

1.3

Vitamin A and vitamin C

differ in the colors displayed

on their electrostatic poten-

tial maps. The areas of blue

and red in vitamin C’s map

indicate differences in the

electronegativity of the

atoms in this molecule.

These differences result

in the water solubility of

vitamin C.

7.9 Reactivity 289

side of each period possess an increased nuclear charge. Both of these factors

cause electrons to be attracted to the nucleus more strongly.

The decrease in electronegativity that we note as we move down groups

makes sense because the atoms are growing larger.With an extra intervening shell

of electrons at each step, the nuclear charge is being screened from the electrons

shared in a covalent bond. The effect of the increased size and electron screening

outweighs the increase in nuclear charge as we move down a group.

The trends in electronegativity, electron afﬁnity, and ionization energy are

not without exceptions, so they are general trends rather than absolute rules.

They do, however, reveal one very clear message about which atoms are most

likely to attract electrons and become negative ions, and which are most likely to

lose electrons and become positive ions. Atoms on the extreme left of the periodic

table, especially the lower extreme left, will most readily form positive ions. Atoms on

the extreme right of the periodic table, especially the upper extreme right (always

remembering that Group VIIIA is excluded), will most readily form negative ions.

EXERCISE 7.8 Trends in Electronegativity

Without examining the actual Pauling electronegativity values, predict which of

these elements in each pair is more electronegative.

Li or Be N or P S or I

Solution

Be is closer in the periodic table to ﬂuorine. It is more electronegative.

N is only two elements away from ﬂuorine, whereas P is three elements away.

Nitrogen is more electronegative.

S is two elements away from ﬂuorine. It is more electronegative.

PRACTICE 7.8

Which of the following in each pair is more electronegative?

P or Na Ne or Cl N or C

See Problems 68–70.

7.9 Reactivity

The concept of reactivity is a descriptive but very useful notion in chemistry.

When we describe an element as “highly reactive,” we mean that it readily partic-

ipates in chemical reactions to form compounds. When we describe an element

as “unreactive,” we mean it does not readily participate in chemical reactions to

form compounds. As we have already discovered, most highly reactive elements

are, in their natural state, combined with other elements in the form of com-

pounds. You will never ﬁnd a hunk of soft, shiny sodium in nature. Rather, you

will ﬁnd it in ionic form in seawater or in salts on land. Unreactive elements can

be found in their pure form. This has great bearing on the uses we make of ele-

ments. The metals sodium and potassium, for example, would be of little use for

making automobiles or jewelry, because they react so vigorously and readily with

water. Gold, on the other hand, is an ideal metal for jewelry. Because it is so unre-

active, it can be found in its free form within the earth. The scarcity of gold makes

it unsuitable for the construction of automobiles, even though such automobiles

might last for a very long time.

We compromise when making automobiles, and build their metal compo-

nents mostly out of steel, which is largely composed of the element iron, as we

Gold in its natural state.

290 Chapter 7 Periodic Properties of the Elements

noted in the NanoWorld/MacroWorld feature in Section 7.2. Iron is not totally

unreactive; it corrodes into rust (iron oxide) when combined with the oxygen

found in air or dissolved in water. Iron lasts long enough, however, to form the

basic materials of automobiles that will last quite a few years.

Can the periodic table tell us which elements are the most reactive? The ele-

ments that react most readily generally do so by either losing or gaining electrons.

The most reactive elements will be those with the greatest tendency to form pos-

itive ions (those with the lowest ionization energies), or with the greatest ten-

dency to form negative ions (those with the largest negative electron afﬁnities or

the largest electronegativities). These highly reactive elements are found at the

extreme left of the periodic table (Groups IA and IIA) and on the far right in

Groups VIA and VIIA. The least reactive elements of all are found in Group

VIIIA, the noble gases; and those with intermediate and low reactivities are found

between the extremes just mentioned, in the central portion of the periodic table.

In Group IB, we ﬁnd gold and silver, unreactive elements we use in jewelry, and

also copper, which is sufﬁciently unreactive to be used in copper electric wires

and the pipes of plumbing systems.

Our discussion now enables us to ﬁnd more general degrees of logic in the pe-

riodic table. We now have enough understanding to give at least rudimentary an-

swers to the questions we raised in the chapter opening about why oxygen reacts

with so many elements and why gold has a very low reactivity. These questions

can be answered by considering the location of these elements in the periodic

La*

†

104

105

106

107

108

109

110

Uun

111

Uuu

IIIB

IVB

VIB

VIIB

IIB

100

101

102

103

IIA

IIIA

IVA

VIA

VIIA

VIIIA

*Lanthanides

†

Actinides

112

Uub

Period

VIIIB

810

Periodic table with the most reactive elements highlighted in green.

Application

Oxyhemoglobin Hemoglobin

table and all that this means. Another question posed in our introduction was

what makes iron so well-suited to the task of carrying oxygen around the body,

bound to the iron-containing protein we call hemoglobin. The iron in hemo-

globin is in the form of Fe

ions, which bind reversibly (that is, they can continue

to bind and then release, depending on conditions) to oxygen molecules (O

The binding has to be readily reversible, because hemoglobin must combine with

oxygen where oxygen is abundant, in the blood vessels of the lungs; but must re-

lease the oxygen to the tissues of the body in which oxygen levels are much lower.

A positive ion can perform the binding task by interacting with the electrons of

the oxygen molecule. In order to do that in a readily reversible way, however, the

ion and the atom from which it is derived must have neither too strong nor too

weak an attraction for electrons. A suitable ion will be from an element of inter-

mediate electronegativity, likely to be found around the middle of the periodic

table. Iron, in the middle of the periodic table, can do the job very nicely. This is

only one of several reasons why iron ions serve within hemoglobin molecules as

the oxygen carriers of life, but it is surely a signiﬁcant one. Other reasons include

the availability of iron in the environment in which life originated and the suit-

ability of iron to combining with the other components of hemoglobin.

One idea that we did not discuss here but will consider in the next two chap-

ters is that judgments about the reactivity of an element cannot be made in isola-

tion. That is, atoms react with other types of atoms, and the extent to which a

reaction occurs depends on the properties of both, or even many, types of atoms.

More broadly speaking, the reaction environment must be considered in decid-

ing the nature of chemical behavior.

The Reactivity Series of Metals

The relative reactivities of some of the most abundant and most useful elements

in the environment are very signiﬁcant to us. For example, in deciding which

metal is suitable for a particular industrial use, we ﬁnd that the ease with which it

corrodes (oxidizes in the presence of substances in the environment, often lead-

ing to deterioration in the properties of the metal) will be very signiﬁcant. The

most common form of corrosion is reaction with oxygen, and all forms of corro-

sion are chemical reactions of one kind or another. The most reactive metals will

generally corrode most quickly. Thinking about such issues led to the idea of list-

ing metals in a

reactivity series (or activity series), which ranks selected metals in

order of reactivity (see Table 7.14).

Such a list can be made by exposing metals to a range of substances, such as

hydrochloric acid, water, and air, and observing how readily they react. There are

some problems with this approach. For example, aluminum reacts more quickly

with oxygen than does iron. The aluminum soon becomes coated, however, with

a cohesive layer of aluminum oxide that protects the aluminum beneath from

further corrosion. That is why aluminum cookware is quite long-lived. Iron

7.9 Reactivity 291

The four oxygen molecules in oxyhemo-

globin can be seen near the center of

the model. Hydrogen atoms have been

omitted for clarity.

Application

A Reactivity Series

(Activity Series)

Potassium Most reactive

Sodium

Lithium

Calcium

Magnesium

Aluminum

Zinc

Iron

Tin

Lead

Copper

Mercury

Silver

Gold Least reactive

TABLE 7.14

Video Lesson: The Activity Series

of the Elements

oxide, on the other hand, does not form any protective layer but instead produces

the crumbling, weak structure we know as “rust,” which breaks away to reveal

fresh corrosion-prone metal underneath. Nevertheless, the idea of ranking metals

by reactivity has proved very useful.

An alternative, and in some ways more satisfactory, way of ranking metals by

reactivity is found in the “electrochemical series” that we will discuss in Chapter

19. This is related to the tendency of each metal to give up electrons. Remember

that when metals react, they generally do so by losing outer electrons to form

ions. The most reactive metals are the ones that lose electrons and form positive

ions most readily.

7.10 The Elements and the Environment

The elements of planet Earth are distributed throughout the environment in

places and forms that are a consequence of their physical and chemical proper-

ties, and these physical and chemical properties can be related to the elements’

positions in the periodic table. The

environment in this sense comprises the entire

world around us, including the Earth, its atmosphere, and ourselves.

Geologists describe the structure of the Earth in terms of ﬁve distinct regions:

the core, mantle, crust, hydrosphere, and atmosphere. These are shown diagram-

matically in Figure 7.14. Although nobody has traveled to the Earth’s core or has

even been able to obtain samples, we know its structure from the accumulation of

292 Chapter 7 Periodic Properties of the Elements

Rusty aluminum (left) and rusty iron (right).

FIGURE 7.14

The structure of planet Earth.

Application

HEMICAL

ENCOUNTERS:

The Elements and

the Environment

Core

Mantle

Crust

Hydrosphere

Atmosphere

years of indirect evidence. There is no doubt that the core is composed largely of

iron, mixed with smaller amounts of the metals nickel and cobalt and lighter ele-

ments such as carbon and sulfur. The core is largely molten, so these elements

form a molten alloy (a homogeneous mixture of metals) with some nonmetals

mixed within the alloy. The predominance of iron at the center of the Earth is a

consequence of its high density in relation to most of the other elements that

make up our planet. When the Earth formed, it went through a molten phase,

during which the densest materials were drawn toward the center. The less dense

materials rose outward, in just the same way as a mixture of cooking oil and water

will settle into layers, the dense water beneath the less dense oil.

The mantle is another region of the Earth that has never been directly sam-

pled, although a few different types of volcanic rocks are believed to have been

derived from this region. The evidence of these rocks, combined with other indi-

rect evidence, suggests that the mantle is composed largely of oxides of the

elements iron, magnesium, silicon, calcium, and aluminum. The low-density ele-

ment oxygen, which is a gas at typical atmospheric temperatures and pressures,

has been retained within the Earth’s mantle because of its chemical reactivity.

This has caused it to react with various metals to form the much more dense

compounds found in the mantle.

The Earth’s crust makes up less than 1 percent of the mass of the planet and,

as shown in Figure 7.15, it has a ratio of elements that differs greatly from that of

the Earth as a whole. However, it is the only solid part of the Earth that we can an-

alyze directly. The crust is also the part of the environment from which we draw

nearly all of our raw materials for use in industrial fabrication and as sources of

energy. The most abundant elements of the crust are listed in Table 7.15. These

elements are largely found bound within compounds, oxides being the most pre-

dominant. Oxides of silicon are among the most common components of the

crust. For example, silicon dioxide (SiO

) is the principal component of sand and

of the many types of rock from which sand is derived.

The hydrosphere is the name given to the waters of the Earth, comprising

oceans, seas, lakes, rivers, and underground aquifers. In addition to the hydrogen

and oxygen that make up water, the hydrosphere contains a great variety of dis-

solved substances that inﬂuence its properties and suitability for various uses.

Seawater is unsafe for humans to drink because of its high concentration of salts,

such as sodium chloride and magnesium chloride. Earth’s atmosphere contains

the low-density, gaseous materials that nevertheless are sufﬁciently dense to be

retained by the gravitational pull of the planet. As shown in Table 7.16, the at-

mosphere is largely composed of nitrogen and oxygen, with much smaller

amounts of argon, carbon dioxide, and other rare gases. Some of the most signif-

icant components of the environment, however, are too rare even to show up in

most tables of the atmosphere’s composition. Gases such as sulfur dioxide and

7.10 The Elements and the Environment 293

Oil

Crust Whole earth

Iron

Oxygen

Silicon

Magnesium

Calcium

Aluminum

Sodium

Potassium

Sulfur

Nickel

Other

FIGURE 7.15

Comparison of the elements in the Earth’s crust with those in the entire Earth.

Vinegar

Oil and vinegar, which is

mostly water, are immisci-

ble; they do not mix. The

denser vinegar sinks to

the bottom of this bottle.

nitrogen dioxide are extremely rare overall, but they cause problems over areas

where they are present as environmental pollutants.

Finally, it is worth pointing out that the element hydrogen (H

), the most

abundant element in the universe, is of such low density that any hydrogen re-

leased into the atmosphere will eventually escape into space, unless it reacts with

some other chemicals on the way. The only way for hydrogen to be retained on

our planet is within chemical compounds, such as water (H

O), methane (CH

or any one of millions of other compounds, including most of those found

within living things. If chemical changes had not locked up hydrogen within such

chemical compounds, we would never have evolved to ponder and study these

chemical changes.

294 Chapter 7 Periodic Properties of the Elements

The Abundance of Elements in the Earth’s Crust

Element Abundance (percent, by mass)

Oxygen 46.60

Silicon 27.72

Aluminum 8.13

Iron 5.00

Calcium 3.63

Sodium 2.83

Potassium 2.59

Magnesium 2.09

Titanium 0.44

Hydrogen 0.14

Phosphorus 0.12

Manganese 0.10

Copper 0.0070

Gold 0.00000005

TABLE 7.15

Composition of Dry Air

Substance Volume Fraction

0.7808

0.2095

Ar 0.00934

0.00034

Ne 1.82 × 10

−5

He 5.82 × 10

−6

2 × 10

−6

Kr 1.1 × 10

−6

5 × 10

−7

O5× 10

−7

Xe 8.7 × 10

−8

< 1 × 10

−6

< 1 × 10

−7

,CO,NO,I

< 1 × 10

−8

TABLE 7.16

The Bottom Line

■

The structure of the periodic table was initially de-

veloped by scientists trying to make sense of the dif-

fering reactivities of all the elements found in the

natural environment. (Section 7.1)

■

The structure of the periodic table includes “blocks”

deﬁned in terms of which type of orbital is being

ﬁlled as we imagine ﬁlling up the available orbitals

using the Aufbau principle. This gives us the s-block,

p-block, d-block, and f-block. (Section 7.1)

■

The horizontal “period” of the periodic table to

which an element belongs indicates how many

energy levels are either fully or partially occupied by

electrons in atoms of that element. (Section 7.1)

■

Elements in any one main group (Groups IA

through VIIIA) have the same number of electrons

in their highest energy level. (Section 7.1)

■

The modern form of the table ﬁrst began to take

shape through the work of the German Julius

Lothar Meyer and the Russian Dmitri Mendeleev.

(Section 7.1)

■

The elements in the periodic table are arranged into

three main sections: the metals, nonmetals, and met-

alloids (or semimetals). (Section 7.2)

■

Metals and nonmetals exhibit speciﬁc types of physi-

cal properties. (Section 7.2)

■

The periodic table is divided into groups, and ele-

ments in each group are similar in electronic struc-

ture, physical properties, and chemical reactivity.

(Section 7.3)

■

Periodicity is apparent in the way important physical

and chemical characteristics recur in a periodic

manner as we move through the periodic table. (Sec-

tion 7.4)

■

Examples of characteristics that show clear periodic-

ity are atomic size (Section 7.5), ionization energy

(Section 7.6), electron afﬁnity (Section 7.7), electro-

negativity (Section 7.8), and reactivity (Section 7.9).

■

The distribution of the elements on planet Earth is a

result of their chemical reactivities. (Section 7.10)

Key Words 295

Key Words

actinides The second set of inner transition elements.

The actinides include elements 90–103. (p. 275)

activity series A series of elements ranked in order of re-

activity. Also known as a reactivity series. (p. 291)

alkali metals The Group IA elements, excluding hydro-

gen. (p. 270)

alkaline earth metals The Group IIA elements. (p. 270)

alloy A homogeneous mixture of metals. (p. 266)

atomic radius Half the distance between the nuclei in a

molecule consisting of identical atoms. Also known

as the covalent radius. (p. 280)

catalyst A substance that greatly speeds up the rate of a

reaction without being consumed by that reaction.

(p. 272)

chalcogens The Group VIA elements. (p. 273)

covalent radius Half the distance between the nuclei in a

molecule consisting of identical atoms. Also known

as the atomic radius. (p. 280)

electron afﬁnity The energy change associated with the

addition of an electron to an atom in the gaseous

state. (p. 286)

electronegativity The relative ability of an atom partici-

pating in a chemical bond to attract electrons to

itself. (p. 287)

environment The world around us, comprising the

Earth and its atmosphere. We ourselves are part of

the environment. (p. 292)

ﬁrst ionization energy (I

) The energy required to remove

one electron from each atom of an element in the

gaseous state. (p. 282)

halogens The Group VIIA elements. (p. 274)

inert gases The Group VIIIA elements. Also known as

noble gases. (p. 274)

inner transition elements The elements in the f-block of

the periodic table, consisting of the lanthanides and

actinides. (p. 275)

ionization energy The energy needed to ionize an atom

(or molecule or ion) by removing an electron from it.

(p. 282)

lanthanides The ﬁrst set of inner transition elements.

The lanthanides include elements 58–71. (p. 275)

main groups Groups IA through VIIIA of the periodic

table. Groups IB through VIIIB are not main groups.

(p. 269)

metallic radius Half the distance between the nuclei

of the atoms within the solid structure of a metal.

(p. 280)

metalloids A small number of elements that have char-

acteristics midway between those of the metals and

those of the nonmetals. Also known as semimetals.

(p. 268)

metals A large number of elements that share certain

typical characteristics, including shiny appearance,

malleability, and ability to conduct electricity. (p. 265)

noble gases The Group VIIIA elements. Also known as

inert gases. (p. 274)

nonmetals A collection of elements that shares certain

characteristics that are in contrast to those of the

metals, including being gases or dull, brittle solids at

room temperature. (p. 268)

octet Eight electrons in the valence shell of an atom.

(p. 275)

periodicity The recurrence of characteristic properties as

we move through a series, such as the elements of the

periodic table. (p. 279)

reactivity The ability of an element or compound to

participate in chemical reactions. Reactivity is an im-

precise but useful term, particularly when we are dis-

cussing relative reactivities under deﬁned conditions.

(p. 289)

reactivity series A series of elements ranked in order of

reactivity. Also known as an activity series. (p. 291)

second ionization energy (I

) The energy required to

remove one electron from each singly charged +1 ion

of an element in the gaseous state. (p. 282)

semimetals A small number of elements that have char-

acteristics midway between those of the metals and

those of the nonmetals. Also known as metalloids.

(p. 268)

transition elements Metal elements from Groups IB

through VIIIB found in the middle of the periodic

table. (p. 275)

valence electrons The electrons that occupy the outer-

most energy level of an atom, which interact with

those of other atoms. (p. 269)

The answers to the odd-numbered problems and some selected

problems appear at the back of the book, as represented by the

blue numbering.

Section 7.1 The Big Picture—Building the Periodic Table

Skill Review

1. Cite two examples wherein two elements’ locations on the

current periodic table would be reversed if they were placed

in accordance with Mendeleev’s atomic mass system of

organization.

2. Some may argue that providing a list of elements in alpha-

betical order along with their properties would be the most

useful arrangement of the known elements. Explain what ad-

vantage there is to arranging them in the periodic manner

displayed in our traditional periodic table format.

3. Identify the “block” location in the periodic table for each of

these elements.

a. Sr c. S e. Se g. Sb

b. Sc d. Sn f. Sm

4. Identify the “block” location in the periodic table for each of

these elements.

a. Rb c. Ru e. Re g. Ra

b. Rn d. Rh f. Rf

5. Identify the “block” location in the periodic table for each of

these elements.

a.C c.Cu e.Cs g.Ce

b. Ca d. Cr f. Cl

6. Identify the “block” location in the periodic table for each of

these elements.

a. He c. Fe e. Xe g. Te

b. Be d. Ge f. Ne

Chemical Applications and Practices

7. The following portion of a hypothetical periodic table shows

the approximate melting points of elements that are neigh-

bors to a missing element. On the basis of the data provided,

predict the melting point of the missing element.

2610°C

3000°C ? 3180°C

8. The following portion of a hypothetical periodic table shows

the densities of elements that are neighbors to a missing ele-

ment. On the basis of the data provided, predict the density

of the missing element.

10.2 g/cm

16.7 g/cm

? 21.0 g/cm

9. In each of these pairs, select the larger of the two elements.

a. Mg or Ca c. Cl or Br

b. P or S d. Cs or Ba

10. In each of these pairs, select the larger of the two elements.

a. K or Ca c. Se or Br

b. O or S d. Ra or Ba

Section 7.2 The First Level of Structure—Metals,

Nonmetals, and Metalloids

Skill Review

11. Classify each of these as a metal, nonmetal, or metalloid.

a. Zn c. Ge e. Si g. Se

b. Ga d. Sn f. P h. As

12. Classify each of these as a metal, nonmetal, or metalloid.

a. Li c. Co e. C g. I

b. Ne d. Te f. Bi h. Sb

13. Most elements, at room temperature, are either solids or

gases. However, two elements, at room temperature, are clas-

siﬁed as liquids. What are these two elements?

14. Most of the key ingredients to alloy with iron to make various

types of steel are metals. However, this is not always the case.

What are two nonmetals that can be alloyed with steel?

Chemical Applications and Practices

15. a. In the production of stainless steel, small amounts of sev-

eral elements are alloyed with iron. Two of the chief ingre-

dients are nickel and chromium. Use the chart given in

Table 7.3 to report the number of grams of each metal

found in 100.0 g of stainless steel.

b. The atoms of the added metals can become part of the

iron structure, which changes the properties of iron. How

many atoms of nickel and chromium, respectively, are pre-

sent in the 100.0 g of stainless steel?

16. Assume that the other elements (besides chromium, nickel,

and iron) found in stainless steel are insigniﬁcant. Based on

the calculations in Problem 15, how many iron atoms are

present in the sample for every atom of chromium?

17. Locate the element antimony in the periodic table. Would

this element be classiﬁed as a metal, nonmetal, or metalloid?

Antimony is used in alloys with lead or tin, is combined with

other substances to make ﬁreproof fabrics, is mixed with

some types of glass and ceramics, and serves as the active

component of some medicines. Antimony, however, does

have some toxic properties. One way to obtain it is as follows:

+ Fe → FeS +Sb

Balance this equation. The annual worldwide production of

antimony is approximately 6.00

× 10

tons. How many

moles of antimony is this?

18. Locate the element gold in the periodic table. Would this ele-

ment be classiﬁed as a metal, nonmetal, or metalloid? Gold is

used not only in jewelry but also as a wire in some computer

parts and in other situations where a noncorrosive metal is

needed. Gold can be solubilized in water by using a mixture

of very strong acids, as follows:

Au(s) +HCl(aq) + NO

−

(aq) →

[AuCl

]

−

(aq) + H

O(l) + NO(g)

Balance this equation. How many grams of gold must be

solubilized to produce 1 mol of NO(g)?

296 Chapter 7 Periodic Properties of the Elements

Focus Your Learning

Section 7.3 The Next Level of Structure—Groups in the

Periodic Table

Skill Review

19. Using the most common oxidation number for each, write

out the formulas of the compounds that would be most likely

to form when Li,Be,B,C,N,O,and F,respectively,combine

with oxygen.

20. Write out the chemical formulas for Li, Na, K, Rb, and Cs, re-

spectively, combining with oxygen. What accounts for the

similarities in the formulas of these compounds?

21. If an element from Group VIA were going to form an ionic

compound, which type of element (metal or nonmetal)

would be most likely to form the other part of the com-

pound? Explain your answer.

22. If an element from Group IIA were going to form an ionic

compound, which type of element (metal or nonmetal)

would be most likely to form the other part of the com-

pound? Explain your answer.

23. How many unpaired electrons are found in the elements that

make up Group VIIA? How many unpaired electrons are

found in the most common ions of those elements?

24. How many unpaired electrons are found in the elements that

make up Group IIIA? How many unpaired electrons are

found in the most common ions of those elements?

Chemical Applications and Practices

25. The listing in Table 7.13 of elements in the human body is

based on mass. Considering only the ﬁrst three (oxygen,

carbon, and hydrogen), indicate whether the ranking would

change if the table were based on number of atoms instead of

mass. Explain the basis of your answer.

26. The listing in Table 7.13 of elements in the human body

includes calcium. What use does our body make of this

element? In addition, calculate the quantity of calcium in a

typical human in units of milligrams per kilogram of body

weight.

27. Many dietary supplements contain a list of metals present.

This list may include zinc, iron, cobalt, and others. Name

one important role that transition metals play in human

biochemistry.

28. Iron plays several roles in humans. One of its important roles

involves the electron transport chain. Iron can easily change

oxidation states between the +2 and +3 charges. What char-

acteristic feature of iron’s electron conﬁguration makes this

possible?

Section 7.4 The Concept of Periodicity

Skill Review

29. The noble gases, except for helium, have a stable, ﬁlled octet.

Explain why helium is still considered a noble gas even

though it does not have a stable, ﬁlled octet.

30. The alkali metals have a complete shell of a noble gas plus an

extra electron. Explain why hydrogen is often associated with

this group of the periodic table.

Chemical Applications and Practices

31. Examine the full electron conﬁguration of calcium and

potassium. Explain why calcium is not as reactive as

potassium.

32. The same reasoning used to answer Problem 31 applies to the

relative reactivity of magnesium compared to sodium, but

why is sodium not as reactive as potassium?

33. In some of the original periodic table research, tellurium was

placed after iodine due to its slightly greater relative mass.

However, this placed tellurium in Group VIIA directly below

bromine, chlorine, and ﬂuorine. After examining the elec-

tron conﬁguration of tellurium and iodine explain why the

present classiﬁcation is logical.

34. Nickel is placed on the periodic table immediately after

cobalt. However, its average mass is less than that of cobalt.

Explain why the placement is logical.

Section 7.5 Atomic Size

Skill Review

35. These three spheres can be used to compare the relative sizes

of sodium, magnesium, and sulfur. Which sphere would best

represent each atom?

36. These three spheres can be used to compare the relative sizes

of potassium, rubidium, and cesium. Which sphere would

best represent each atom?

37. If the bond length for a COC bond were 154 pm, what would

you determine as the radius of a carbon atom?

38. If the bond length for a HOH bond were 75 pm, what would

you determine as the approximate radius for an atom of

hydrogen?

39. If the bond length for a COCl bond were 171 pm, what

would you determine as the approximate radius for an atom

of chlorine? (Use the answer from Problem 37 to assist you in

determining the radius of a chlorine atom.)

40. If the bond length for a HOF bond were 92 pm, what would

you determine as the approximate radius for an atom of ﬂu-

orine? (Use the answer from Problem 38 to assist you in de-

termining the radius of a ﬂuorine atom.)

H–H bond

75 pm

C–C bond

154 pm

Focus Your Learning 297