6.1 Organizing the Elements 6t1lara.weebly.com/uploads/1/6/3/2/1632178/ch6pdf.pdf · The Periodic...

The Periodic Table 155

Print• Guided Reading and Study Workbook,

Section 6.1• Core Teaching Resources, Section 6.1,

Review, Interpreting Graphics• Transparencies, T65–T66

Technology• Interactive Textbook with ChemASAP,

Assessment 6.1• Go Online, Section 6.1

6.1

FOCUSObjectives6.1.1 Explain how elements are

organized in a periodic table.6.1.2 Compare early and modern

periodic tables.6.1.3 Identify three broad classes of

elements.

Guide for Reading

Build VocabularyWord Parts Point out that metal, non-metal, and metalloid all have the same root. Discuss the meanings of the suffix -oid (resembling) and prefix non- (not).

Reading StrategyUse Prior Knowledge Ask students to write down three things they know about the development of the periodic table. After they read the section, have them revise their statements and record three new things that they learned.

INSTRUCT

Have students look at the opening photograph and describe the types of characteristics that might be used to group produce in a store. Explain that an initial classification may be broad, as in produce dairy, and canned goods. Then these categories can be subdi-vided.

Searching For an Organizing Principle

TEACHER DemoTEACHER Demo

Organizing Elements

Display as many samples of elements as possible. If elements are not avail-able, use photographs. Ask students to use properties of the elements to orga-nize them into groups, and to provide an explanation for their groupings.

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Section Resources

Connecting to Your World

Section 6.1 Organizing the Elements 155

Chlorine35.453 amu

Bromine79.904 amu

Iodine126.90 amu

6.1 Organizing the Elements

In 1916, a self-service grocery store opened in Memphis, Tennessee. Shoppers could select items from shelves instead of waiting for a clerk to gather the items for them. In a self-

service store, the customers must know how to find the products. From your experience, you know

that products are grouped according to sim-ilar characteristics. You don’t expect to find fresh fruit with canned fruit, or bottled

juice with frozen juice. With a logical clas-sification system, finding and comparing

products is easy. In this section, you will learn how elements are arranged in the

periodic table and what that arrangement reveals about the elements.

Guide for Reading

Key Concepts• How did chemists begin to

organize the known elements?• How did Mendeleev organize

his periodic table?• How is the modern periodic

table organized?• What are three broad classes

of elements?

Vocabularyperiodic law

metals

nonmetals

metalloid

Reading StrategyComparing and Contrasting As you read, compare and contrast Figures 6.4 and 6.5. How are these two versions of the periodic table similar? How are they different?

Searching For an Organizing PrincipleA few elements have been known for thousands of years, including copper,silver, and gold. Yet only 13 elements had been identified by the year 1700.Chemists suspected that other elements existed. They had even assignednames to some of these elements, but they were unable to isolate the ele-ments from their compounds. As chemists began to use scientific methodsto search for elements, the rate of discovery increased. In one decade(1765–1775), chemists identified five new elements, including three color-less gases—hydrogen, nitrogen, and oxygen. Was there a limit to the num-ber of elements? How would chemists know when they had discovered allthe elements? To begin to answer these questions, chemists needed to finda logical way to organize the elements.

Chemists used the properties of elements to sort them into groups.In 1829, a German chemist, J. W. Dobereiner (1780–1849), published a clas-sification system. In his system, elements were grouped into triads. A triadis a set of three elements with similar properties. The elements in Figure 6.1formed one triad. Chlorine, bromine, and iodine may look different.But they have very similar chemical properties. For example, they reacteasily with metals. Dobereiner noted a pattern in his triads. One elementin each triad tended to have properties with values that fell midwaybetween those of the other two elements. For example, the average ofthe atomic masses of chlorine and iodine is [(35.453 � 126.90)/2] or81.177 amu. This value is close to the atomic mass of bromine, whichis 79.904 amu. Unfortunately, all the known elements could not begrouped into triads.

Figure 6.1 Chlorine, bromine, and iodine have very similar chemical properties. The numbers shown are the average atomic masses for these elements.

156 Chapter 6

Section 6.1 (continued)

Mendeleev's Periodic TableUse VisualsFigure 6.3 Have students study Men-deleev's version of the periodic table. Show students how the elements with similar chemical properties were arranged in horizontal rows. Point out the gaps in the table where Mendeleev left room for elements that had not yet been discovered. Emphasize that Men-deleev did not make the first periodic table, but he did make the first table that gained wide acceptance.

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156 Chapter 6

Mendeleev’s Periodic TableFrom 1829 to 1869, different systems were proposed, but none of themgained wide acceptance. In 1869, a Russian chemist and teacher, DmitriMendeleev, published a table of the elements. Later that year, a Germanchemist, Lothar Meyer, published a nearly identical table. Mendeleev wasgiven more credit than Meyer because he published his table first andbecause he was better able to explain its usefulness. The stamp in Figure 6.2is one of many ways that Mendeleev’s work has been honored.

Mendeleev developed his table while working on a textbook for his stu-dents. He needed a way to show the relationships among more than 60 ele-ments. He wrote the properties of each element on a separate note card.This approach allowed him to move the cards around until he found anorganization that worked. The organization he chose was a periodictable. Elements in a periodic table are arranged into groups based on aset of repeating properties. Mendeleev arranged the elements in hisperiodic table in order of increasing atomic mass.

Figure 6.3 is an early version of Mendeleev’s periodic table. Look at thecolumn that starts with Ti � 50. Notice the two question marks betweenthe entries for zinc (Zn) and arsenic (As). Mendeleev left these spaces in histable because he knew that bromine belonged with chlorine and iodine. Hepredicted that elements would be discovered to fill those spaces, and hepredicted what their properties would be based on their locations in thetable. The elements between zinc and arsenic were gallium and germa-nium, which were discovered in 1875 and 1886, respectively. There was aclose match between the predicted properties and the actual properties ofthese elements. This match helped convince scientists that Mendeleev’speriodic table was a powerful tool.

Figure 6.2 Dimitri Mendeleev proposed a periodic table that could be used to predict the properties of undiscovered elements.

Figure 6.3 In this early version of Mendeleev’s periodic table, the rows contain elements with similar properties.Observing A fourth element is grouped with chlorine (Cl), bromine (Br), and (I) iodine. What is this element’s symbol?


The Periodic LawUse VisualsFigure 6.4 This is the first of three periodic tables in the chapter. The information in this first table is limited so that students can see how all the elements fit in the table, and can focus on atomic number as an organizing principle. Have students compare Figures 6.3 and 6.4. Ask, How is the modern periodic table shown here similar to the early periodic table? (In both tables, elements are arranged in rows and columns. Early periodic tables ordered the elements according to atomic mass. Modern periodic tables order elements according to atomic number.) What does an element’s position in the table indicate about its properties? (The properties of an element are similar to those of other ele-ments in the same group.)

DiscussTo emphasize the concept that ele-ments in the same group of the peri-odic table have similar properties, point out one group and describe some properties of the elements in that group. You can point to the group on the far right, and explain that all of these elements are gases that do not normally form compounds. Using a large poster-size display of the periodic table, trace one period across the table. Describe some physical and chemical properties of each element in the period. Then have students compare these elements to those below them in the next period.

Word OriginsOne meaning of perimeter is “a line or boundary around an area.” It comes from peri, “around,” and meter, “measure.”

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Answers to...Figure 6.3 FFigure 6.4 8

Checkpoint

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Word Origins

1

H1

H

3

Li

11

Na

37

Rb

55

Cs

87

Fr

19

K

2

He

10

Ne

18

Ar

54

Xe

86

Rn

36

Kr

9

F

17

Cl

53

I

85

At

35

Br

8

O

16

S

52

Te

84

Po

34

Se

7

N

15

P

51

Sb

83

Bi

33

As

6

C

14

Si

50

Sn

82

Pb

114

Uuq

111

Uuu

112

Uub

108

Hs109

Mt110

Ds

32

Ge

5

B

13

Al

49

In

81

Tl

31

Ga

48

Cd

80

Hg

30

Zn

47

Ag

79

Au

29

Cu

46

Pd

78

Pt

28

Ni

45

Rh

77

Ir

27

Co

44

Ru

76

Os

26

Fe

107

Bh

43

Tc

75

Re

25

Mn

106

Sg

42

Mo

74

W

24

Cr

105

Db

41

Nb

73

Ta

23

V

104

Rf

40

Zr

72

Hf

22

Ti

103

Lr

39

Y

71

Lu

21

Sc

102

No

70

Yb

101

Md

69

Tm

100

Fm

68

Er

99

Es

67

Ho

98

Cf

66

Dy

97

Bk

65

Tb

96

Cm

64

Gd

95

Am

63

Eu

94

Pu

62

Sm

93

Np

61

Pm

92

U

60

Nd

91

Pa

59

Pr

90

Th

58

Ce

89

Ac

57

La

4

Be

12

Mg

38

Sr

56

Ba

88

Ra

20

Ca

1

2

3

4

5

6

7

The Periodic LawThe atomic mass of iodine (I) is 126.90. The atomic mass of tellurium (Te)is 127.60. Based on its chemical properties, iodine belongs in a group withbromine and chlorine. So Mendeleev broke his rule and placed telluriumbefore iodine in his periodic table. He assumed that the atomic masses foriodine and tellurium were incorrect, but they were not. Iodine has a smalleratomic mass than tellurium does. A similar problem occurred with otherpairs of elements. The problem wasn’t with the atomic masses but withusing atomic mass to organize the periodic table.

Mendeleev developed his table before scientists knew about the struc-ture of atoms. He didn’t know that the atoms of each element contain aunique number of protons. Remember that the number of protons is theatomic number. In 1913, a British physicist, Henry Moseley, determined anatomic number for each known element. Tellurium’s atomic number is52 and iodine’s is 53. So it makes sense for iodine to come after telluriumin the periodic table. In the modern periodic table, elements arearranged in order of increasing atomic number.

The elements in Figure 6.4 are arranged in order of atomic number,starting with hydrogen, which has atomic number 1. There are seven rows,or periods, in the table. Period 1 has 2 elements, Period 2 has 8 elements,Period 4 has 18 elements, and Period 6 has 32 elements. Each period corre-sponds to a principal energy level. There are more elements in higher num-bered periods because there are more orbitals in higher energy levels.(Recall the rules you studied in Chapter 5 for how electrons fill orbitals.)

The elements within a column, or group, in the periodic table havesimilar properties. The properties of the elements within a period changeas you move across a period from left to right. However, the pattern ofproperties within a period repeats as you move from one period to the next.This pattern gives rise to the periodic law: When elements are arranged inorder of increasing atomic number, there is a periodic repetition of theirphysical and chemical properties.

Checkpoint How many periods are there in a periodic table?

Figure 6.4 In the modern periodic table, the elements are arranged in order of increasing atomic number. Interpreting Diagrams How many elements arethere in the second period?

Periodic comes from the Greek roots peri meaning “around” and hodos, mean-ing “path.” In a periodic table, properties repeat from left to right across each period. The Greek word metron means “measure.”What does perimetermean?

158 Chapter 6


Metals, Nonmetals, and MetalloidsUse VisualsFigure 6.5 The focus in Figure 6.5 is on the classification of elements as metals, nonmetals, and metalloids. Point out the sections of the table that correspond to metals, nonmetals, and metalloids. Emphasize the periodic trend across each period from metals to nonmetals. Ask, Which class of ele-ments do most of the elements belong to? (metals) What do the three numbers at the top of each col-umn represent? (different systems for numbering the groups)

Download a worksheet on Metals and Nonmetals for students to complete, and find additional teacher support from NSTA SciLinks.

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Differentiated Instruction

158 Chapter 6

Metals, Nonmetals, and MetalloidsMost periodic tables are laid out like the one in Figure 6.5. Some elementsfrom Periods 6 and 7 are placed beneath the table. This arrangement makesthe periodic table more compact. It also reflects an underlying structure ofthe periodic table, which you will study in Section 6.2. Each group in thetable in Figure 6.5 has three labels. Scientists in the United States used thelabels shown in red. Scientists in Europe used the labels shown in blue.There is some overlap between the systems, but in many cases two differ-ent groups have the same letter and number combination.

For scientists to communicate clearly, they need to agree on the stan-dards they will use. The International Union of Pure and Applied Chemis-try (IUPAC) is an organization that sets standards for chemistry. In 1985,IUPAC proposed a new system for labeling groups in the periodic table.They numbered the groups from left to right 1 through 18 (the black labelsin Figure 6.5). The large periodic table in Figure 6.9 includes the IUPAC sys-tem and the system used in the United States. The latter system will bemost useful when you study how compounds form in Chapters 7 and 8.

Dividing the elements into groups is not the only way to classify thembased on their properties. The elements can be grouped into three broadclasses based on their general properties. Three classes of elements aremetals, nonmetals, and metalloids. Across a period, the properties of ele-ments become less metallic and more nonmetallic.

Metals The number of yellow squares in Figure 6.5 shows that most ele-ments are metals—about 80 percent. Metals are good conductors of heatand electric current. A freshly cleaned or cut surface of a metal will have ahigh luster, or sheen. The sheen is caused by the metal’s ability to reflectlight. All metals are solids at room temperature, except for mercury (Hg).Many metals are ductile, meaning that they can be drawn into wires. Mostmetals are malleable, meaning that they can be hammered into thin sheetswithout breaking. Figure 6.6 shows how the properties of metals can deter-mine how metals are used.

For: Links on Metals and Nonmetals

Visit: www.SciLinks.orgWeb Code: cdn-1061

MetalloidsMetals Nonmetals

2IIA

2A

3IIIA

3B

4IVA

4B

5VA

5B

6VIA

6B

7VIIA

7B

8 11IB

1B

12IIB

2B

13IIIB3A

14IVB4A

15VB5A

16IVB6A

17VIB7A

18VIIB8A

10

1IA1A

9VIIIA

8B

1

H

3

Li11

Na

37

Rb55

Cs87

Fr

19

K

2

He10

Ne18

Ar

54

Xe86

Rn

36

Kr

9

F17

Cl

53

I85

At

35

Br

8

O16

S

52

Te84

Po

34

Se

7

N15

P

51

Sb83

Bi

33

As

6

C14

Si

50

Sn82

Pb114

Uuq111

Uuu112

Uub108

Hs109

Mt110

Ds

32

Ge

5

B13

Al

49

In81

Tl

31

Ga48

Cd80

Hg

30

Zn47

Ag79

Au

29

Cu46

Pd78

Pt

28

Ni45

Rh77

Ir

27

Co44

Ru76

Os

26

Fe

107

Bh

43

Tc75

Re

25

Mn

106

Sg

42

Mo74

W

24

Cr

105

Db

41

Nb73

Ta

23

V

104

Rf

40

Zr72

Hf

22

Ti

103

Lr

39

Y71

Lu

21

Sc

102

No

70

Yb101

Md

69

Tm100

Fm

68

Er99

Es

67

Ho98

Cf

66

Dy97

Bk

65

Tb96

Cm

64

Gd95

Am

63

Eu94

Pu

62

Sm93

Np

61

Pm92

U

60

Nd91

Pa

59

Pr90

Th

58

Ce89

Ac

57

La

4

Be12

Mg

38

Sr56

Ba88

Ra

20

Ca

Figure 6.5 One way to classify elements in the periodic table is as metals, nonmetals, and metalloids. Inferring What is the purpose for the black stair-step line?

Less Proficient ReadersMake a chart on the board of the three types of elements. Ask students to fill in details about the general characteristics of the three classes as they read about them.

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CLASS ActivityCLASS

Name The ElementHave each student choose an element without revealing the choice to other students. Each student should write a short description of the chosen ele-ment to read to the class. Ask the other students to identify the element from its description. Encourage students to be as specific as possible in their descriptions. Example: bromine is a reddish-brown liquid.

FYIAt a height of 630 feet, The Gateway Arch is the tallest monument in the United States. It commemorates the westward expansion of the U. S. in the nineteenth century.

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Answers to...Figure 6.5 Students are likely to say the line separates the metals from the nonmetals.

Checkpoint

metals


Iron (Fe)The Gateway Arch in St. Louis, Missouri, is covered in stainless steel containing iron and two other metals, chromium (Cr) and nickel (Ni). The steel is shiny, malleable, and strong. It also resists rusting.

Copper (Cu)Copper is ductile and second to only silver as a conductor of electric current. The copper used in electrical cables must be 99.99% pure.

Aluminum (Al)Aluminum is one of the metals that can be shaped into a thin sheet, or foil. To qualify as a foil, a metal must be no thicker than about 0.15 mm.

Nonmetals In Figure 6.5, blue is used to identify the nonmetals. Theseelements are in the upper-right corner of the periodic table. There is agreater variation in physical properties among nonmetals than amongmetals. Most nonmetals are gases at room temperature, including the maincomponents of air—nitrogen and oxygen. A few are solids, such as sulfurand phosphorus. One nonmetal, bromine, is a dark-red liquid.

The variation among nonmetals makes it difficult to describe one set ofgeneral properties that will apply to all nonmetals. However, nonmetals arenot metals, as their name implies. So they tend to have properties that areopposite to those of metals. In general, nonmetals are poor conductors ofheat and electric current. Carbon is an exception to this rule. Solid nonmet-als tend to be brittle, meaning that they will shatter if hit with a hammer.

Checkpoint Which type of elements tend to be good conductors of heat and electric current?

Figure 6.6 The metals iron, copper, and aluminum have many important uses. How each metal is used is determined by its properties.

160 Chapter 6


DiscussPoint out silicon's position in the peri-odic table and explain that silicon is a metalloid, an element having both metallic and nonmetallic properties. Ask students about the properties that distinguish metals from nonmetals. Have them name other elements that are classified as metalloids. (Chemists do not always agree on which ele-ments to classify as metalloids. Some include polonium (Po) and astatine (At) as metalloids. Some classify polonium as a metal and astatine as a nonmetal.)

ASSESSEvaluate UnderstandingHave students draw a concept map relating the following terms: groups, periods, periodic law, periodic table, repeating properties, metals, nonmetals, and metalloids.

ReteachCompare and contrast the periodic table and a monthly calendar. Similarity: The same progression of days occurs each week. Difference: Although the length of a day is based on a natural event (Earth’s rotation), what happens on a Monday or Saturday depends on human decisions, not on an underlying natural principle.

Connecting Concepts

The atomic number of an element tells you how many protons are in its nucleus. It is a good way to organize elements because atomic number is unique for each element, while other properties, such as atomic mass, can vary for atoms of an element.

with ChemASAP

If your class subscribes to the Interactive Textbook, use it to review key concepts in Section 6.1.

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Section 6.1 Assessment1. Chemists used the properties of elements

to sort them into groups.2. in order of increasing atomic mass3. in order of increasing atomic number4. metal, metalloid, nonmetal5. b

6. a. metal b. metalloid c. nonmetal d. metal

7. lithium, potassium, rubidium, cesium, or francium

160 Chapter 6

withChemASAP

Metalloids There is a heavy stair-step line in Figure 6.5 that separates themetals from the nonmetals. Most of the elements that border this line areshaded green. These elements are metalloids. A metalloid generally hasproperties that are similar to those of metals and nonmetals. Under someconditions, a metalloid may behave like a metal. Under other conditions,it may behave like a nonmetal. The behavior often can be controlled bychanging the conditions. For example, pure silicon is a poor conductorof electric current, like most nonmetals. But if a small amount of boron ismixed with silicon, the mixture is a good conductor of electric current, likemost metals. Silicon can be cut into wafers, like those being inspectedin Figure 6.7, and used to make computer chips.

6.1 Section Assessment

1. Key Concept How did chemists begin the process of organizing elements?

2. Key Concept What property did Mendeleev use to organize his periodic table?

3. Key Concept How are elements arranged in the modern periodic table?

4. Key Concept Name the three broad classes of elements.

5. Which of these sets of elements have similar physical and chemical properties?

a. oxygen, nitrogen, carbon, boron b. strontium, magnesium, calcium, beryllium c. nitrogen, neon, nickel, niobium

6. Identify each element as a metal, metalloid, or nonmetal.

a. gold b. silicon c. sulfur d. barium

7. Name two elements that have properties similar to those of the element sodium.

Atomic Number What does an atomic number tell you about the atoms of an element? Why is atomic number better than atomic mass for organizing the elements in a periodic table? Use what you learned in Section 4.2 to answer this question.

Assessment 6.1 Test yourself on the concepts in Section 6.1.

Figure 6.7 Pancake-sizedcircular slices of silicon, called wafers, are used to make computer chips. Because a tiny speck of dust can ruin a wafer, the people who handle the wafers must wear “bunny” suits. The suits prevent skin, hair, or lint from clothing from entering the room’s atmosphere.



Section 6.2• Core Teaching Resources,

Section 6.2 Review• Transparencies, T67–T69• Laboratory Manual, Lab 9


Assessment 6.2• Go Online, Section 6.2

6.2

FOCUSObjectives6.2.1 Describe the information in a

periodic table.6.2.2 Classify elements based on

electron configuration.6.2.3 Distinguish representative ele-

ments and transition metals.

Guide for Reading

Build VocabularyLINCS Have students use the LINCS strategy. In LINCS exercises, students List what they know about each term, Imagine a picture that describes the term, Note a reminding “sound-alike” word, Connect the terms to the sound-alike word by making up a short story, and then perform a brief Self-test.

Reading StrategyPredict After students preview the sec-tion, have them predict the relationship between an element’s electron configu-ration and its position in the periodic table. Students can revise their predic-tions after they read the section.

INSTRUCT

After students read the opening para-graph, have them volunteer other exam-ples of small items that contain lots of information (e.g., train schedules, digital watch face, nutrient labels).

Squares in the Periodic TableUse VisualsFigure 6.8 Remind students what aver-age atomic mass represents. Make sure students understand the vertical column of numbers.

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Answers to...Figure 6.8 There are 11 protons in the nucleus and 11 electrons in the three occupied energy levels.

Section Resources

Connecting to Your World

Section 6.2 Classifying the Elements 161

6.2 Classifying the Elements

The sculptor Augustus Saint-Gaudens designed this gold coin at the request of Theodore Roosevelt. President Roosevelt wanted coins minted in the United States to be as beautiful as ancient Greek coins, which he admired. The coin is an example

of a double eagle. The name derives from the fact that the coin was worth twice as much as $10

coins called eagles. A coin may contain a lot of information in a small space—its value, the year it was minted, and its country of origin. Each square in a peri-

odic table also contains a lot of informa-tion. In this section, you will learn what

types of information are usually listed in a periodic table.

Guide for Reading

Key Concepts• What type of information

can be displayed in a periodic table?

• How can elements be classified based on their electron configurations?

Vocabularyalkali metals

alkaline earth metals

halogens

noble gases

representative elements

transition metal

inner transition metal

Reading StrategyRelating Text and VisualsAs you read, look carefully at Figure 6.9. After you read the section, explain what you can tell about an element from the color assigned to its square and the color assigned to its symbol.

Squares in the Periodic TableThe periodic table displays the symbols and names of the elements,

along with information about the structure of their atoms. Figure 6.8shows one square from the detailed periodic table of the elements in Figure6.9 on page 162. In the center of the square is the symbol for sodium (Na).The atomic number for sodium (11) is above the symbol. The elementname and average atomic mass are below the symbol. There is also a verti-cal column with the numbers 2, 8, and 1, which are the number of electronsin each occupied energy level of a sodium atom.

The symbol for sodium is printed in black because sodium is a solid atroom temperature. In Figure 6.9, the symbols for gases are in red. The sym-bols for the two elements that are liquids at room temperature, mercuryand bromine, are in blue. The symbols for some elements in Figure 6.9 areprinted in green. These elements are not found in nature. In Chapter 25,you will learn how scientists produce these elements.

The background colors in the squares are used to distinguish groups ofelements. For example, two shades of gold are used for the metals inGroups 1A and 2A. The Group 1A elements are called alkali metals, and theGroup 2A elements are called alkaline earth metals. The name alkali comesfrom the Arabic al aqali, meaning “the ashes.” Wood ashes are rich in com-pounds of the alkali metals sodium and potassium. Some groups of non-metals also have special names. The nonmetals of Group 7A are calledhalogens. The name halogen comes from the combination of the Greekword hals, meaning salt, and the Latin word genesis, meaning “to be born.”There is a general class of compounds called salts, which include the com-pound called table salt. Chlorine, bromine and iodine, the most commonhalogens, can be prepared from their salts.

Figure 6.8 This is the element square for sodium from the periodic table in Figure 6.9. Interpreting Diagrams Whatdoes the data in the square tell you about the structure of sodium atoms?

281

11

NaSodium22.990

Atomic number

Electrons in eachenergy level

Element symbol

Element name

Averageatomic mass

162 Chapter 6


Use VisualsFigure 6.9 Have students examine the periodic table and find the types of information that the table provides. Summarize suggestions on the board. Examples include element name, atomic number, average atomic mass, and physical state at atmospheric pres-sure and room temperature.

DiscussHave students examine the periodic table and determine the accuracy of this statement: Atomic mass always increases as atomic number increases. (The trend is generally true, but there are exceptions. For example, the atomic num-bers of Co, Ni, and Cu increase by one unit (27, 28, 29), but Ni has the lowest average atomic mass of these three elements.)

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Facts and FiguresElement NomenclatureNames suggested by those who create new elements must be approved by the Interna-tional Union of Pure and Applied Chemistry (IUPAC). The process can be lengthy. For example, the names chosen for elements 104–108 by the nomenclature committee of IUPAC in 1994 were not those endorsed by the American Chemical Society (ACS). After

years of negotiation, a compromise resulted in the names that appear in the periodic table. Elements beyond 111 have been discovered but are not yet named. Element 110 was discovered in 1994 by a team of scientists in Darmstadt, Germany. The IUPAC Council voted formal approval of the name darmstadtium (Ds) on August 16, 2003.

162 Chapter 6

SolidC

LiquidBrBr

GasHe

Not foundin natureTcTc

Alkali Metals

Metalloids

Other Metals

Alkaline Earth Metals

Transition Metals

Inner transition metals

Nonmetals

Noble Gases

Transition ElementsRepresentative Elements

1

HHydrogen

1.0079

103

LrLawrencium

(262

3

LiLithium

6.941

19

KPotassium

39.098

37

RbRubidium

85.468

4

BeBeryllium

9.0122

12

MgMagnesium

24.305

11

NaSodium

22.990

20

CaCalcium

40.08

38

SrStrontium

87.62

56

BaBarium

137.33

88

RaRadium

(226)

71

LuLutetium

174.97

72

HfHafnium

178.49

73

TaTantalum

180.95

74

WTungsten

183.85

75

ReRhenium

186.21

76

OsOsmium

190.2

104

RfRutherfordium

(261)

105

DbDubnium

(262)

106

SgSeaborgium

(263)

107

BhBohrium

(264)

108

HsHassium

(265)

109

MtMeitnerium

(268)

77

IrIridium

192.22

39

YYttrium

88.906

40

ZrZirconium

91.22

41

NbNiobium

92.906

42

MoMolybdenum

95.94

43

TcTechnetium

(98)

44

RuRuthenium

101.07

45

RhRhodium

102.91

21

ScScandium

44.956

22

TiTitanium

47.90

23

VVanadium

50.941

24

CrChromium

51.996

25

MnManganese

54.938

26

FeIron

55.847

27

CoCobalt

58.933

55

CsCesium

132.91

87

FrFrancium

(223)

57

LaLanthanum

138.91

89

AcActinium

(227)

90

ThThorium

232.04

92

UUranium

238.03

93

NpNeptunium

(237)

94

PuPlutonium

(244)

91

PaProtactinium

231.04

58

CeCerium

140.12

60

NdNeodymium

144.24

59

PrPraseodymium

140.91

61

PmPromethium

(145)

62

SmSamarium

150.4

Periodic Table of the Elements

2

2A

3

3B4

4B5

5B6

6B7

7B8 9

8B

1

1A1

21

22

281

282

2881

2882

28

1881

28

1882

28

1818

81

28

1818

82

28

183218

81

28

183218

82

28

183232

92

2892

28

1892

28

1832

92

28

102

28

1810

2

28

183210

2

28

18323210

2

28

112

28

1812

1

28

183211

2

28

18323211

2

28

131

28

1813

1

28

183212

2

28

18323212

2

28

132

28

1814

1

28

183213

2

28

18323213

2

28

142

28

1815

1

28

183214

2

28

18323214

2

28

152

28

18161

28

1832152

28

183232152

28

182482

28

1823

82

28

1822

82

28

1821

82

28

1820

82

28

1818

92

28

183218

92

28

18321810

2

28

183220

92

28

18322192

28

183222

92

28

18322482

Lanthanide Series

Actinide Series

Figure 6.9 In this periodic table, the colors of the boxes are used to classify representative elements and transition elements.


RelateHave students start a list of elements found in the body that are essential for metabolism. Explain that many ele-ments are present in trace amounts but are, nevertheless, essential for sur-vival. Have students do research on such elements as copper, chromium, iodine, and manganese to find out more about their biological roles. They can begin with the discussion of micro-nutrients on page R45 of the Elements Handbook.

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28

NiNickel

58.71

29

CuCopper

63.546

30

ZnZinc

65.38

31

GaGallium

69.72

32

GeGermanium

72.59

33

AsArsenic

74.922

34

SeSelenium

78.96

35

BrBromine

79.904

36

KrKrypton

83.80

51

SbAntimony

121.75

52

TeTellurium

127.60

53

IIodine

126.90

54

XeXenon

131.30

81

TlThallium

204.37

49

InIndium

114.82

82

PbLead

207.2

83

BiBismuth

208.98

84

PoPolonium

(209)

85

AtAstatine

(210)

86

RnRadon

(222)

50

SnTin

118.69

13

AlAluminum

26.982

14

SiSilicon

28.086

15

PPhosphorus

30.974

16

SSulfur

32.06

17

ClChlorine

35.453

18

ArArgon

39.948

5

BBoron

10.81

6

CCarbon

12.011

7

NNitrogen

14.007

8

OOxygen

15.999

9

FFluorine

18.998

10

NeNeon

20.179

2

HeHelium

4.0026

46

PdPalladium

106.4

47

AgSilver

107.87

48

CdCadmium

112.41

78

PtPlatinum

195.09

110

DsDarmstadtium

(269)

111

UuuUnununium

(272)

112

UubUnunbium

(277)

114

UuqUnunquadium

79

AuGold

196.97

80

HgMercury

200.59

63

EuEuropium

151.96

64

GdGadolinium

157.25

65

TbTerbium

158.93

66

DyDysprosium

162.50

67

HoHolmium

164.93

68

ErErbium

167.26

69

TmThulium

168.93

70

YbYtterbium

173.04

95

AmAmericium

(243)

96

CmCurium

(247)

97

BkBerkelium

(247)

98

CfCalifornium

(251)

99

EsEinsteinium

(252)

100

FmFermium

(257)

101

MdMendelevium

(258)

102

NoNobelium

(259)

14

SiSilicon

28.086

284

11

1B12

2B

13

3A14

4A15

5A16

6A17

7A

18

8A

10

2

28

162

28

1818

28

1832171

28

183232171

28

182582

28

18322582

28

181

28

18181

28

1832181

28

183232181

28

182592

28

18322592

28

182

28

183218

3

28

1818

3

28

183

283

23

28

183218

4

28

1818

4

28

184

284

24

28

183218

5

28

1818

5

28

185

285

25

28

183218

6

28

1818

6

28

186

286

26

28

183218

7

28

1818

7

28

187

287

27

28

183218

8

28

1818

8

28

188

288

28

28

18182

28

1832182

28

183232182

28

182782

28

18322782

28

1828

82

28

183228

82

28

1829

82

28

183229

82

28

1830

82

28

183230

82

28

1831

82

28

183231

82

28

1832

82

28

183232

82

* * *

*Name not officially assigned.

Atomic number

Electrons in eachenergy level

Element symbol

Element name

Averageatomic mass

164 Chapter 6


Electron Configurations in GroupsDiscussLead a class discussion on electron configurations of noble gases and representative elements. Select some elements and have students write out the electron configurations for those elements. Have students compare the electron configurations for all the elements in a single group. Ask stu-dents to identify similarities. (Noble gases are sometimes classified as rep-resentative elements because they are in the p block of elements.)

FYIThe term valence electron will be intro-duced in Chapter 7.

Download a worksheet on Chemi-cal Families for students to com-plete, and find additional teacher support from NSTA SciLinks.

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Using Shorthand NotationYou can use the electron configurations of noble gases to produce a shorthand notation for other electron configurations. The symbol of a noble gas in brackets can stand for its electron configuration. The configuration of sodium can be written as [Ne] 3s1; the config-uration of fluorine as [He] 2s22p5; and the configuration of chromium as [Ar] 3d54s1.

164 Chapter 6

Figure 6.10 This blimp contains helium, one of the noble gases.Applying Concepts Whatdoes the ability of a helium-filled blimp to rise in air tell you about the density of helium?

For: Links on Chemical Families


Electron Configurations in GroupsElectrons play a key role in determining the properties of elements. Sothere should be a connection between an element’s electron configurationand its location in the periodic table. Elements can be sorted into noblegases, representative elements, transition metals, or inner transition metalsbased on their electron configurations. You may want to refer to Figure 6.9as you read about these classes of elements.

The Noble Gases The blimp in Figure 6.10 is filled with helium. Heliumis an example of a noble gas. The noble gases are the elements in Group 8Aof the periodic table. These nonmetals are sometimes called the inert gasesbecause they rarely take part in a reaction. The electron configurations forthe first four noble gases in Group 8A are listed below.

Look at the description of the highest occupied energy level for each ele-ment, which is highlighted in yellow. The s and p sublevels are completelyfilled with electrons. Chapter 7 will explain how this arrangement of elec-trons is related to the relative inactivity of the noble gases.

The Representative Elements Figure 6.11 shows the portion of theperiodic table containing Groups 1A through 7A. Elements in these groupsare often referred to as representative elements because they display a widerange of physical and chemical properties. Some are metals, some are non-metals, and some are metalloids. Most of them are solids, but a few aregases at room temperature, and one, bromine, is a liquid.

In atoms of representative elements, the s and p sublevels of the high-est occupied energy level are not filled. Look at the electron configurationsfor lithium, sodium, and potassium. In atoms of these Group 1A elements,there is only one electron in the highest occupied energy level. The electronis in an s sublevel.

In atoms of carbon, silicon, and germanium, in Group 4A, there are fourelectrons in the highest occupied energy level.

For any representative element, its group number equals the number ofelectrons in the highest occupied energy level.

Checkpoint Why are noble gases sometimes referred to as inert gases?

Helium (He) 1s2

Neon (Ne) 1s22s22p6

Argon (Ar) 1s22s22p63s23p6

Krypton (Kr) 1s22s22p63s23p63d104s24p6

Lithium (Li) 1s22s1

Sodium (Na) 1s22s22p63s1

Potassium (K) 1s22s22p63s23p64s1

Carbon (C) 1s22s22p2

Silicon (Si) 1s22s22p63s23p2

Germanium (Ge) 1s22s22p63s23p63d104s24p2

Facts and Figures



Differences in Reactivity of Metals Purpose Students observe differences in the reactivity of magnesium, tin, and copper.

Materials 0.2M HCl; 6 large test tubes; test tube rack; small pieces of clean magnesium, tin, and copper; matches; wood splint

Safety Wear goggles for this demo.

Procedure Place three large test tubes in a test tube rack. To each test tube, add a small, clean piece of a dif-ferent metal—magnesium, tin, and copper. Carefully add some of the 0.2M HCl to each test tube and invert a test tube over it. Point out the appearance of bubbles. After a while, carefully ignite the hydrogen in the inverted test tube from the magnesium or tin reaction.

Expected Outcome The magnesium and tin react with the HCl, producing hydrogen. The hydrogen ignites with an explosive pop. The copper shows no sign of reaction.

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Answers to...Figure 6.10 Helium is less dense than air.

Checkpoint

They rarely take part in a reaction.


1

HHydrogen

1.0079

3

LiLithium

6.941

19

KPotassium

39.098

37

RbRubidium

85.468

4

BeBeryllium

9.0122

12

MgMagnesium

24.305

11

NaSodium22.990

20

CaCalcium

40.08

38

SrStrontium

87.62

56

BaBarium137.33

88

RaRadium

(226)

31

GaGallium

69.72

32

GeGermanium

72.59

33

AsArsenic74.922

34

SeSelenium

78.96

35

BrBromine79.904

51

SbAntimony

121.75

52

TeTellurium

127.60

53

IIodine126.90

81

TlThallium204.37

49

InIndium114.82

82

PbLead207.2

83

BiBismuth208.98

84

PoPolonium

(209)

85

AtAstatine

(210)

50

SnTin

118.69

13

AlAluminum

26.982

14

SiSilicon28.086

15

PPhosphorus

30.974

16

SSulfur32.06

17

ClChlorine35.453

5

BBoron10.81

6

CCarbon12.011

7

NNitrogen

14.007

8

OOxygen15.999

9

FFluorine18.998

55

CsCesium132.91

87

FrFrancium

(223)

1

21

22

281

282

2881

2882

28

1881

28

1882

28

181881

28

181882

28

18321881

28

18321882

28

1832183

28

18183

28

183

283

23

28

1832184

28

1818

4

28

184

284

24

28

183218

5

28

1818

5

28

185

285

25

28

183218

6

28

1818

6

28

186

286

26

28

1832187

28

18187

28

187

287

27

1A

2A 3A 4A 5A 6A 7A

Figure 6.11 Some of the representative elements exist in nature as elements. Others are found only in compounds.

Sodium When salt lakes evaporate, they form salt pans like this one in Death Valley, California. The main salt in a salt pan is sodium chloride.

Magnesium This magnified view of a leaf shows the green structures where light energy is changed into chemical energy. The compound chlorophyll, which contains magnesium, absorbs the light.

Arsenic This bright red ore is a major source of arsenic in Earth’s crust. It contains a compound of arsenic and sulfur.

Sulfur These scientists are sampling gases being released from a volcano through a vent called a fumarole. The yellow substance is sulfur.

166 Chapter 6


Transition Elements

CLASS ActivityCLASS

Lanthanides in Consumer ProductsHave students use the library or Inter-net to compile a list of consumer prod-ucts that contain lanthanides or require lanthanides for processing.

DiscussStudents may think that elements whose names are familiar are always more abundant than less familiar ele-ments. You can use the following data to show why calling inner transition metals “rare-earth elements” is mis-leading. In Earth’s crust, there is 1.7 ppb of osmium, 37 ppb of platinum, 67 ppb of mercury, 1700 ppb of tantalum, 6000 ppb of thorium, 10,000 ppb of lead, 60,000 ppb of cerium, and 34,000 ppb of lanthanum.

Use VisualsFigure 6.12 Emphasize the relation-ship between the position of an ele-ment in the periodic table and the element’s atomic structure. Ask, What do all the elements in Group 1A have in common? (They all have one electron in the s orbital in the highest occupied energy level.) What do all the ele-ments in Group 3A have in common? (They all have one electron in a p orbital in the highest occupied energy level.) To help students understand the instruc-tions for writing electron configura-tions for elements in the d and f blocks, review the aufbau diagram in Figure 5.7 on p. 133.

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Differentiated InstructionLess Proficient ReadersDivide the class into small groups. Have each group write electron configurations for an s-block element, a p-block element, a d-block element, and an f-block element.

Don’t assign any of the d-block elements with upredictable configurations. Ask stu-dents to explain the method they used to determine specific electron configurations.

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166 Chapter 6

Transition ElementsIn the periodic table, the B groups separate the A groups on the left side ofthe table from the A groups on the right side. Elements in the B groups,which provide a connection between the two sets of representative ele-ments, are referred to as transition elements. There are two types of transi-tion elements—transition metals and inner transition metals. They areclassified based on their electron configurations.

The transition metals are the Group B elements that are usually dis-played in the main body of a periodic table. Copper, silver, gold, and ironare transition metals. In atoms of a transition metal, the highest occupied ssublevel and a nearby d sublevel contain electrons. These elements arecharacterized by the presence of electrons in d orbitals.

The inner transition metals appear below the main body of the peri-odic table. In atoms of an inner transition metal, the highest occupied ssublevel and a nearby f sublevel generally contain electrons. The innertransition metals are characterized by f orbitals that contain electrons.Before scientists knew much about inner transition metals, people beganto refer to them as rare-earth elements. This name is misleading becausesome inner transition metals are more abundant than other elements.

Blocks of Elements If you consider both the electron configurationsand the positions of the elements in the periodic table, another patternemerges. In Figure 6.12, the periodic table is divided into sections, orblocks, that correspond to the highest occupied sublevels. The s block con-tains the elements in Groups 1A and 2A and the noble gas helium. The pblock contains the elements in Groups 3A, 4A, 5A, 6A, 7A, and 8A, with theexception of helium. The transition metals belong to the d block, and theinner transition metals belong to the f block.

You can use Figure 6.12 to help determine electron configurations ofelements. Each period on the periodic table corresponds to a principalenergy level. Say an element is located in period 3. You know that the s andp sublevels in energy levels 1 and 2 are filled with electrons. You read acrossperiod 3 from left to right to complete the configuration. For transition ele-ments, electrons are added to a d sublevel with a principal energy level thatis one less than the period number. For the inner transition metals, theprincipal energy level of the f sublevel is two less than the period number.This procedure gives the correct electron configurations for most atoms.

Figure 6.12 This diagram classifies elements into blocks according to sublevels that are filled or filling with electrons.Interpreting Diagrams In the highest occupied energy level of a halogen atom, how many electrons are in the p sublevel?

s block

p block

d block

f block

d5 d6 d7 d8d1 d2 d3 d4 d9 d10

f 5 f 6 f7 f 8f1 f 2 f 3 f 4 f 9 f10f11 f12f13f14

p5 p6p1 p2 p3 p4

s2

s1

s2


CONCEPTUAL PROBLEM 6.1

Answers8. a. 1s22s22p2

b. 1s22s22p63s23p63d104s24p65s2

c. 1s22s22p63s23p63d34s2

9. a. B, Al, Ga, In, Tlb. F, Cl, Br, I, At c. Ti, Zr, Hf, Rf

Practice Problems PlusChapter Review problems 34 and 35 are related to Conceptual Problem 6.1.

ASSESSEvaluate UnderstandingCall out pairs of elements in the same group and have students write their electron configurations. This activity can be made into a game where groups compete with other groups to come up with the answer first. Students should eventually be able to write the electron configurations quickly.

ReteachTo reinforce the relationship between configurations and position in the peri-odic table, provide configurations and ask students to identify and locate the elements. Ask students to explain which parts of a configuration proved most useful for determining the identity.

Elements Handbook

Electric current passes through a tube filled with gas, and the gas emits light: pale yellow for helium, red orange for neon, lavender for argon, gray for krypton, blue for xenon. The color depends on the electron configuration.

with ChemASAP


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Answers to...Figure 6.12 five

Section 6.2 Assessment10. symbols and names of the elements;

atomic number and average atomic mass; information about electron configuration

11. representative elements, noble gases, transition metals, and inner transition metals

12. They are in the same group and have the same number of electrons in the highest occupied energy level.

13. a. noble gas b. transition metal c. representative element

14. Cu, Cd, Au, Co15. 5


Handbook

withChemASAP

CONCEPTUAL PROBLEM 6.1

Using Energy Sublevels to Write Electron ConfigurationsNitrogen is an element that organisms need to remain healthy. How-ever, most organisms cannot obtain nitrogen directly from air. A few organisms can convert elemental nitrogen into a form that can be absorbed by plant and animal cells. These include bacteria that live in lumps called nodules on the roots of legumes. The photograph shows the nodules on a bean plant. Use Figure 6.12 to write the elec-tron configuration for nitrogen (N), which has atomic number 7.

Solve Apply concepts to this situation.

Nitrogen is located in the second period of the periodic table and in the third group of the pblock. Nitrogen has seven electrons. Based on Figure 6.12, the configuration for the two elec-trons in the first energy level is 1s 2. The config-uration for the five electrons in the second energy level is 2s 22p 3.

8. Use Figure 6.9 and Figure 6.12 to write the elec-tron configurations of the following elements.a. carbon b. strontium c. vanadium

(Hint: Remember that the principal energy level number for elements in the d block is always one less than the period number.)

9. List the symbols for all the elements whose electron configurations end as follows. Each nrepresents an energy level.a. ns 2np1

b. ns 2np5

c. ns 2nd 2


10. Key Concept What information can be included in a periodic table?

11. Key Concept Into what four classes can elements be sorted based on their electron configurations?

12. Why do the elements potassium and sodium have similar chemical properties?

13. Classify each element as a representative element, transition metal, or noble gas.

a. 1s22s22p63s23p63d104s24p6

b. 1s22s22p63s23p63d 64s2

c. 1s22s22p63s23p2

14. Which of the following elements are transition metals: Cu, Sr, Cd, Au, Al, Ge, Co?

15. How many electrons are in the highest occupied energy level of a Group 5A element?

Noble Gases Look at the atomic properties of noble gases on page R36. In a gas discharge tube, what color light is produced by each noble gas? Use what you know about the structure of atoms to explain why the color is different for each gas.


Practice Problems

Analyze Identify the relevant concepts.

For all elements, the atomic number is equal to the total number of electrons. For a representa-tive element, the highest occupied energy level is the same as the number of the period in which the element is located. From the group in which the element is located, you can tell how many electrons are in this energy level.

168 Chapter 6

True Colors

PurposeThe connection of this Technology & Society to Chapter 6 content may not be immediately obvious. It presents an important application of transition metals. Many pigments depend on the tendency of most transition metals to form colored compounds.

BackgroundPaint is used to cover or hide a surface to which it is applied, decorate a sur-face, or protect a surface. Egg yolk (in egg tempera paint), gum Arabic (in water colors and gouache), and linseed oil (oil paint) are examples of binders. Many pigments depend on the ten-dency of transition metals to form col-ored compounds.

The bison painting is from the Altamira Caves in Spain (about 12,000 B.C.) Although the oldest known paintings are about 35,000 years old, archaeolo-gists in Zambia found pigments and paint grinding equipment that were between 350,000 and 400,000 years old. Some prehistoric artists chemically altered pigments before applying them to cave walls. About one quarter of the samples from Troubat Cave in the Pyrenees were heated in an open fire.

168 Chapter 6

Yellow ochre

Red ochre

Charred wood is asource of charcoal.

True Colors

Paint consists essentially of a pigment, a binder, and a liquid

in which the other components are dissolved or dispersed.

The liquid keeps the mixture thin enough to flow. The

binder attaches the paint to the surface being painted, and

the pigment determines the color. Pigments may be natural

or manufactured. They may be inorganic or organic. The

same pigment can be used in a water-based or oil-based

paint. Comparing and Contrasting Describe at least three

differences between the cave painting and the painting by

Jacob Lawrence.

Natural pigments A prehistoric artist had a limited choice of colors— black from charcoal and red, brown, and yellow from oxides of iron in Earth’s crust. These oxides (or ochre) pigments are often referred to as earth tones.

Prehistoric art Around 14,000 years ago, an artist painted this bison on the ceiling of a cave in Spain. It is about two meters long.


Cobalt BlueManganese Violet

ChromiumOxide Green

Zinc White

Cobalt Yellow

Cadmium OrangeRed Iron Oxide

Section 6 169

From pigments to paintArtists mixed manufactured pigments with binders and solvents to make paint. Although premixed paints became available around 1800, some artists, including Jacob Lawrence, continued to mix their own paints.

Manufactured pigmentsAlchemists (and the Chemists) made pigments that don’t exist in nature. They also made purer versions of natural pigments. Many of these pigments contain transition metals.

The Builders, 1974, by Jacob Lawrence

Coding Scheme for PaintsBecause pigments with the same com-mon name can contain different sub-stances, the Society of Dyers and Colourists in London and the American Association of Textile Chemists and Colorists established a standard coding scheme for pigments. Manufacturers agreed to list both common pigment names and color index codes on labels. See the table below for the composi-tion and codes for the pigments shown on p. 169.

Red iron oxide iron(III) oxide, Fe2O3 PR 101

Cadmium orange solid solution of cadmium selenide, CdSe, and cadmium sulfate, CdSO4

PO 20

Cobalt yellow potassium cobalt nitrite, K3[Co(NO2)6]⋅H2O PY 40

Zinc white zinc oxide, ZnO PW 4

Chromium oxide green chromiu(III) oxide, Cr2O3 PG 17

Cobalt blue cobalt(II) aluminate, Co(AlO2)2 PB 28

Manganese violet manganese ammonium pyrophosphate, Mn(III)NH4P2O7 PV 16

Answers to...Comparing and Contrasting Possible differences include the sub-ject matter, the range of colors, and the anonymity of the artist who painted in the cave.

170 Chapter 6


Section 6.3• Core Teaching Resources,

Section 6.3 Review• Transparencies, T70–T74• Small-Scale Chemistry Laboratory

Manual, Lab 9• Laboratory Practicals, 6-1, 6-2


Animation 7, Assessment 6.3• Go Online, Section 6.3

6.3

FOCUSObjectives6.3.1 Describe trends among the

elements for atomic size.6.3.2 Explain how ions form.6.3.3 Describe periodic trends for

first ionization energy, ionic size, and electronegativity.

Guide for Reading

Build VocabularyGraphic Organizer Have students make a compare/contrast table for cation and anion, and use the table to decide which type of ion an element is likely to form.

Reading StrategyRelate Text and Visuals Tell students that they will learn about trends related to the location of elements in the periodic table. They should look for the visuals throughout the section that summarize the trends.

INSTRUCT

Have students read the opening para-graph. Ask, What does the geometric pattern indicate about the arrange-ment of particles in salt? (It is orderly.)

Trends in Atomic SizeUse VisualsFigure 6.13 Ask, Why can't a scien-tist measure the diameter of a single atom? (because an atom does not have a sharply defined border) Discuss how measuring the distance between nuclei solves this problem.

FYIIn Chapter 8 there are formal defini-tions of molecule and diatomic mol-ecule. The operational definition of a molecule should be sufficient for a discussion of atomic radii.

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Section Resources

170 Chapter 6

6.3 Periodic Trends

An atom doesn’t have a sharply defined boundary. So the radius of an atom cannot be measured directly. There are ways to estimate the sizes of atoms. In one method, a solid is bombarded with X rays, and the paths of the X rays are recorded on film. Sodium chloride (table salt) produced the geometric pattern in the photograph. Such a pattern can be used to calculate the position of nuclei in a solid. The distances between nuclei in a solid are an indication of the size of the parti-cles in the solid. In this section, you will learn how properties such as atomic size are related to the location of elements in the periodic table.

Guide for Reading

Key Concepts• What are the trends among the

elements for atomic size?• How do ions form?• What are the trends among the

elements for first ionization energy, ionic size, and electronegativity?

• What is the underlying cause of periodic trends?

Vocabularyatomic radius

ion

cation

anion

ionization energy

electronegativity

Reading StrategyBuilding Vocabulary After you read this section, explain the difference between a cation and an anion.

Trends in Atomic SizeAnother way to think about atomic size is to look at the units that formwhen atoms of the same element are attached to one another. These unitsare called molecules. Figure 6.13 shows models of molecules (molecularmodels) for seven nonmetals. Because the atoms in each molecule areidentical, the distance between the nuclei of these atoms can be used toestimate the size of the atoms. This size is expressed as an atomic radius.The atomic radius is one half of the distance between the nuclei of twoatoms of the same element when the atoms are joined.

The distances between atoms in a molecule are extremely small. So theatomic radius is often measured in picometers. Recall that there are onetrillion, or 1012, picometers in a meter. Look at the diagram of an iodinemolecule in Figure 6.13. The distance between the nuclei in the molecule is280 pm. Because the atomic radius is one half the distance between thenuclei, a value of 140 pm (280/2) is assigned as the radius of the iodineatom. In general, atomic size increases from top to bottom within agroup and decreases from left to right across a period.

Figure 6.13 This diagram lists the atomic radii of seven nonmetals. An atomic radius is half the distance between the nuclei of two atoms of the same element when the atoms are joined.

Hydrogen (H2)30 pm

Bromine (Br2)120 pm

Iodine (I2)140 pm

Chlorine (Cl2)102 pm

Fluorine (F2)62 pm

Oxygen (O2)68 pm

Nitrogen (N2)70 pm

Distance between nuclei

Atomic radius

Nucleus


Interpreting Graphsa. potassiumb. It increases.c. smallerBecause of the amount of data in this graph, you may need to help students get oriented before they begin to inter-pret the graph.

Enrichment QuestionEmphasize the key roles electrical attraction and repulsion play within atoms and ions. Review the effects of increasing nuclear charge and changes in the shielding effect of electrons on the size of an atom. (Nuclear charge increases within groups and across periods; the shielding effect increases within groups, but is constant across periods.) Have students use these effects to describe the trends for atomic size within a period and within groups.

DiscussAs an analogy to positions and trends in properties of elements in the periodic table, use seating charts and pricing data from local theaters or sports venues to discover trends. Ask students to determine patterns that relate the position of a seat to its price. Students should discover that variables such as distance from the stage or field, location relative to the center of the action, and whether the view will be obstructed, all affect price.

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Answers to...Figure 6.15 the alkali metal

Elements and the Big BangAt the time of the Big Bang, the temperature was many billions of degrees. Neutrons, pro-tons, and electrons may have formed within 10–4 second after the Big Bang, and the light-est nuclei formed within 3 minutes. Matter was in the form of plasma, a sea of positive

nuclei and negative electrons. It took an esti-mated 500,000 years for electrons and nuclei to cool enough to form atoms. According to the Big Bang theory, Earth, with its wealth of chemical elements, formed from the debris of supernova explosions.

Facts and Figures

Section 6.3 Periodic Trends 171

Group Trends in Atomic Size In the Figure 6.14 graph, atomic radius isplotted versus atomic number. Look at the data for the alkali metalsand noble gases. The atomic radius within these groups increases as theatomic number increases. This increase is an example of a trend.

As the atomic number increases within a group, the charge on thenucleus increases and the number of occupied energy levels increases.These variables affect atomic size in opposite ways. The increase in positivecharge draws electrons closer to the nucleus. The increase in the number ofoccupied orbitals shields electrons in the highest occupied energy levelfrom the attraction of protons in the nucleus. The shielding effect is greaterthan the effect of the increase in nuclear charge. So the atomic size increases.

Periodic Trends in Atomic Size Look again at Figure 6.14. In general,atomic size decreases across a period from left to right. Each element hasone more proton and one more electron than the preceding element.Across a period, the electrons are added to the same principal energy level.The shielding effect is constant for all the elements in a period. The increas-ing nuclear charge pulls the electrons in the highest occupied energy levelcloser to the nucleus and the atomic size decreases. Figure 6.15 summa-rizes the group and period trends in atomic size.

Atomic Radius Versus Atomic Number

50

0 10 20 30

Atomic number

40 50 60

100

150

Ato

mic

rad

ius (

pm

)

200

250

300

Sc

Zn

Cd

Perio

d 1

Period 2Period 3

Period 4 Period 5

HeHeNeNe

LiLi

NaNa

KK

RbRbCsCs

KrKr

XeXe

ArAr

Figure 6.14 This graph plots atomic radius versus atomic number for 55 elements.

INTERPRETING GRAPHS

Figure 6.15 The size of atoms tends to decrease from left to right across a period and increase from top to bottom within a group. Predicting If a halogen and an alkali metal are in the same period, which one will have the larger radius?

a. Analyzing Data Whichalkali metal has an atomic radius of 238 pm?b. Drawing ConclusionsBased on the data for alkali metals and noble gases, how does atomic size change within a group?c. Predicting Is an atom of barium, atomic number 56, smaller or larger than an atom of cesium (Cs)?

Siz

e g

ener

ally

incr

ease

s

Size generally decreases

Trends in Atomic Size

172 Chapter 6


Ions

CLASS ActivityCLASS

Listing ElementsGive students a list of elements. Ask them to locate each element in the periodic table, and decide whether its atoms are likely to form positive or negative ions. Have students make a list of elements that are likely to form positive ions and another list of ele-ments that are likely to form negative ions.

FYIThe major discussion of ions and ionic bonding is in Chapter 7. This section presents the limited information that you will need to discuss the trends in ionization energy.

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172 Chapter 6

withChemASAP

IonsSome compounds are composed of particles called ions. An ion is an atomor group of atoms that has a positive or negative charge. An atom is electri-cally neutral because it has equal numbers of protons and electrons. Forexample, an atom of sodium (Na) has 11 positively charged protons and 11negatively charged electrons. The net charge on a sodium atom is zero[(11�) � (�11) � 0].

Positive and negative ions form when electrons are transferredbetween atoms. Atoms of metallic elements, such as sodium, tend to formions by losing one or more electrons from their highest occupied energylevels. A sodium atom tends to lose one electron. Figure 6.16 compares theatomic structure of a sodium atom and a sodium ion. In the sodium ion,the number of electrons (10) is no longer equal to the number of protons(11). Because there are more positively charged protons than negativelycharged electrons, the sodium ion has a net positive charge. An ion with apositive charge is called a cation. The charge for a cation is written as anumber followed by a plus sign. If the charge is 1�, the number 1 is usuallyomitted from the complete symbol for the ion. So Na� is equivalent to Na1�.

Atoms of nonmetallic elements, such as chlorine, tend to form ions bygaining one or more electrons. A chlorine atom tends to gain one electron.Figure 6.16 compares the atomic structure of a chlorine atom and a chlo-ride ion. In a chloride ion, the number of electrons (18) is no longer equal tothe number of protons (17). Because there are more negatively chargedelectrons than positively charged protons, the chloride ion has a net nega-tive charge. An ion with a negative charge is called an anion. The charge foran anion is written as a number followed by a minus sign.

Checkpoint What is the difference between a cation and an anion?

Figure 6.16 When a sodium atom loses an electron, it becomes a positively charged ion. When a chlorine atom gains an electron, it becomes a negatively charged ion.Interpreting Diagrams Whathappens to the protons and neutrons during this change?

Animation 7 Discover the ways that atoms of elements combine to form compounds.

Lose one electron� 1e�

Sodium atom (Na) Sodium ion (Na�)

Nucleus11 p�

12 n0

Nucleus11 p�

12 n0

11 e�

10 e�

Chlorine atom (Cl) Chloride ion (Cl�)

Gain one electron �1e�

Nucleus17 p�

18 n0

Nucleus17 p�

18 n0

17 e�

18 e�


Trends in Ionization EnergyDiscussExplain that ionization energy is a mea-sure of the difficulty in removing an electron from the highest occupied energy level. Ask, Why is the first ion-ization energy of a nonmetal much higher than that of an alkali metal? (Because the nuclear charge increases from left to right across a period and the shielding effect stays the same, it is more difficult to remove an electron.) There are ionization energy graphs in the Ele-ments Handbook on R7, R11, R15, R19, R25, R29, and R33.


Effective Nuclear Charge and Electron ShieldingTo help students understand the con-cepts of effective nuclear charge and electron shielding, choose four students to be “protons” and four to be “electrons.” Construct a lithium “nucleus” by having three protons stand together at the front of the room. Note that for pur-poses of this demo you are ignoring the neutrons. Place two electrons together at a short distance from the nucleus to represent the 1s electrons. Place the third electron a bit farther away to rep-resent the 2s electron. You should be able to draw a line from the nucleus through the 1s electrons to the 2s elec-tron. Point out that there are no other electrons between the 1s electrons and the nucleus. Thus, these electrons expe-rience the full impact of the 3+ charge. Because the third electron’s “view” of the nucleus is partially blocked, it is shielded somewhat from the full force of the 3+ charge.

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Answers to...Figure 6.16 nothing

Checkpoint

A cation has a positive charge. An anion has a neg-ative charge.

FYIBecause the unit for ionization energy is kJ/mol, a footnote in Table 6.1 supplies an operational definition of mole, which is introduced in Sec-tion 10.1.


Table 6.1

Trends in Ionization EnergyRecall that electrons can move to higher energy levels when atoms absorbenergy. Sometimes there is enough energy to overcome the attraction of theprotons in the nucleus. The energy required to remove an electron from anatom is called ionization energy. This energy is measured when an elementis in its gaseous state. The energy required to remove the first electron froman atom is called the first ionization energy. The cation produced has a 1�charge. First ionization energy tends to decrease from top to bottomwithin a group and increase from left to right across a period.

Table 6.1 lists the first, second, and third ionization energies for thefirst 20 elements. The second ionization energy is the energy required toremove an electron from an ion with a 1� charge. The ion produced has a2� charge. The third ionization energy is the energy required to remove anelectron from an ion with a 2� charge. The ion produced has a 3� charge.

Ionization energy can help you predict what ions elements will form.Look at the data in Table 6.1 for lithium (Li), sodium (Na), and potassium(K). The increase in energy between the first and second ionization ener-gies is large. It is relatively easy to remove one electron from a Group 1Ametal atom, but it is difficult to remove a second electron. So Group 1Ametals tend to form ions with a 1� charge.

Ionization Energies of First 20 Elements (kJ/mol *)

*An amount of matter equal to the atomic mass in grams.

Symbol First Second Third

H 1312

He (noble gas) 2372 5247Li 520 7297 11,810Be 899 1757 14,840B 801 2430 3659C 1086 2352 4619N 1402 2857 4577O 1314 3391 5301F 1681 3375 6045Ne (noble gas) 2080 3963 6276Na 496 4565 6912Mg 738 1450 7732Al 578 1816 2744Si 786 1577 3229P 1012 1896 2910S 999 2260 3380Cl 1256 2297 3850Ar (noble gas) 1520 2665 3947K 419 3069 4600Ca 590 1146 4941

174 Chapter 6


Interpreting Graphsa. lithium; sodiumb. First ionization energy decreases as atomic number increases.c. hydrogen; it has only one electron.

Because of the amount of data in this graph, you may need to help students get oriented before they begin to inter-pret the graph.

Enrichment QuestionChallenge students to explain why the portion of the graph for periods 4 and 5 is different from the portion of the graph for periods 2 and 3. (Periods 4 and 5 include transition met-als, whose atoms have electrons in d orbitals.)


Predicting ReactivityPurpose Students observe the rela-tive reactivities of magnesium and cal-cium and predict relative reactivities for other pairs of elements.

Materials 20 mL 1M HCl, two 50-mL beakers, overhead projector, 20 cm magnesium ribbon, 1 g calcium

Safety Wear goggles for this demo.

Procedure Pour 20 mL HCl into each beaker. Set the beakers on an over-head projector. Coil the magnesium ribbon and drop it into one beaker. Drop 1 g calcium into the other beaker. Compare the reaction rates in the two beakers. Point out the positions of the two elements in the periodic table, and relate the difference in reactivity to their first and second ionization ener-gies. Ask students to predict the rela-tive reactivities of other pairs of elements in Groups 1A and 2A.

Expected Outcome The calcium fizzes in the HCl. The magnesium reacts more slowly with the HCl.

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174 Chapter 6

First Ionization Energy Versus Atomic Number

2500

2000

Fir

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izati

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y (

kJ/m

ol)

1500

1000

500

0 10 20 30

Atomic number

40 50 60

He

HBe

N

Ne

LiNa

Mg

PZn As Cd

K RbCs

Kr

Xe

Ar

Group Trends in Ionization Energy Figure 6.17 is a graph of first ion-ization energy versus atomic number. Each red dot represents the data forone element. Look at the data for the noble gases and the alkali metals. Ingeneral, first ionization energy decreases from top to bottom within agroup. Recall that the atomic size increases as the atomic number increaseswithin a group. As the size of the atom increases, nuclear charge has asmaller effect on the electrons in the highest occupied energy level. So lessenergy is required to remove an electron from this energy level and the firstionization energy is lower.

Periodic Trends in Ionization Energy In general, the first ionizationenergy of representative elements tends to increase from left to right acrossa period. This trend can be explained by the nuclear charge, whichincreases, and the shielding effect, which remains constant. The nuclearcharge increases across the period, but the shielding effect remains con-stant. So there is an increase in the attraction of the nucleus for an electron.Thus, it takes more energy to remove an electron from an atom. Figure 6.18summarizes the group and period trends for first ionization energy.

Figure 6.17 This graph reveals group and period trends for ionization energy.

INTERPRETING GRAPHS

a. Analyzing Data Whichelement in period 2 has the lowest first ionization energy? In period 3?b. Drawing ConclusionsWhat is the group trend for first ionization energy for noble gases and alkali metals?c. Predicting If you drew a graph for second ionization energy, which element would you have to omit? Explain.

Figure 6.18 First ionization energy tends to increase from left to right across a period and decrease from top to bottom within a group.Predicting Which element would have the larger first ionization energy—an alkali metal in period 2 or an alkali metal in period 4?

Ener

gy g

ener

ally

dec

reas

es

Energy generally increases

Trends in First Ionization Energy


Quick LABQuick LAB

Periodic Trends in Ionic RadiiObjective After completing this activity, students will be able to• identify periodic trends in ionic size.

Skills Focus Using tables and graphs, predicting, drawing conclusions

Prep Time noneClass Time 40 minutes

Teaching Tips• If time is too limited for students to

make the graph, use Figure 6.19 to answer Questions 1, 2, 4, and 5.

• You may want to reference the radii diagrams in the Elements Handbook on R7, R11, R15, R19, R25, R29, and R33.

Expected OutcomeIonic radii increase from top to bottom within a group. The radii of cations and anions decrease from left to right across a period.

Analyze and Conclude 1. Cations are smaller than their atoms; anions are larger than their atoms.2. Ionic radii increase from top to bot-tom within a group of metals or within a group of nonmetals. 3. Two portions of the curve slope down from left to right. 4. The trend is similar for the periods.5. The radii increase within a group because the number of occupied energy levels increases. The radii of cat-ions decrease across a period because the nuclear charge increases, the shielding effect is constant, and the number of electrons decreases. (The effect is smaller with anions because the number of electrons increases.)

For Enrichment Have students use the graph on page R41 to describe the periodic trend in atomic size for transition metals. Ask how the trend for transition metals compares to the trend for representative elements.

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Answers to...Figure 6.18 an alkali metal in period 2


Quick LABQuick LAB

00

50

100

150

200

2010 30 5040 60Atomic number

Ion

ic r

ad

ii (

pm

)

250

Ionic Radii vs. Atomic Number

Periodic Trends in Ionic Radii

PurposeMake a graph of ionic radius versus atomic number and use the graph to identify periodic and group trends.

Materials• graph paper

ProcedureUse the data presented in Figure 6.19 to plot ionic radius versus atomic number.

Analyze and Conclude1. Describe how the size changes when

an atom forms a cation and when an atom forms an anion.

2. How do the ionic radii vary within a group of metals? How do they vary within a group of nonmetals?

3. Describe the shape of a portion of the graph that corresponds to one period.

4. Is the trend across a period similar or different for periods 2, 3, 4, and 5?

5. Propose explanations for the trends you have described for ionic radii within groups and across periods.

H

Li

K

Rb

Be

MgNa

Ca

Sr

Ba

Ga Ge As Se Br Kr

Sb Te I Xe

Tl

In

Pb Bi Po At Rn

Sn

Al Si P S Cl Ar

B C N O F Ne

He

Cs

30

156

238

255

113

160191

197

215

224

141 122 122 120 120 111

137 139 140 130

172

166

175 170 168 140 140

139

143 109 109 105 102 94

83 77 70 66 62 70

50

273

1+

1+

1+

2+

2+1+

2+

2+

2+

3+ 4+ 3– 2– 1–

5+ 2– 1–

3+

3+

4+ 5+

4+

3+ 4+ 3– 2– 1–

3+ 4+ 3– 2– 1–

1+

60

133

148

44

6695

99

112

134

62 53 222 198 196

62 221 220

95

81

84 74

71

51 41 212 184 181

23 15 146 140 133

169Anion

Cation

Nonmetal

Metalloid

Metal

Atomic radius

Li

156

60 Ionic radius

1A

2A 3A 4A 5A 6A 7A

8A

Figure 6.19 Atomic and ionic radii are an indication of the relative size of atoms and ions. The data listed in Figure 6.19 is reported in picometers (pm).

176 Chapter 6


Trends in Ionic SizeDiscussRelate the periodic trends in ionic size to those discussed earlier for atomic size. Explain that the effective nuclear charge experienced by an electron in the highest occupied orbital of an atom or ion is equal to the total nuclear charge (the number of protons) minus the shielding effect due to electrons in lower energy levels. The effective nuclear charge determines the atomic and ionic radii. As you proceed from left to right in any given period, the principal quantum number, n, of the highest occupied energy level remains constant, but the effective nuclear charge increases. Therefore, atomic and ionic radii decrease as you move to the right in a period. In contrast, within any group, as you proceed from top to bottom, the effective nuclear charge remains nearly constant, but the prin-cipal quantum number, n, increases. Consequently, atomic and ionic radii increase from top to bottom within a group.


Trends in Ionic SizePurpose Students observe an analogy for the effect of adding or removing electrons from an atom.

Materials washers or other small cir-cular items, smaller item to represent the nucleus, overhead projector

Procedure On the overhead projec-tor, make a circle of washers to repre-sent an electron cloud in a neutral atom. The washers should be touching. Place the “nucleus” in the center of the circle. Add or subtract washers to mimic ion formation. With each change, adjust the circle so that the washers are still touching. Explain that the change in the diameter of the circle is analogous to the change in the effec-tive attraction of the nuclear charge for electrons.

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Gifted and TalentedHave students research and describe the phenomenon of the lanthanide contraction. Ask them to discuss how the lanthanide con-traction accounts for the fact that zirconium and hafnium have virtually the same atomic radius even though hafnium is below zirco-nium in Group 4B of the periodic table.

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Differentiated Instruction

176 Chapter 6

Trends in Ionic SizeDuring reactions between metals and nonmetals, metal atoms tend to loseelectrons and nonmetal atoms tend to gain electrons. The transfer has apredictable affect on the size of the ions that form. Cations are alwayssmaller than the atoms from which they form. Anions are always largerthan the atoms from which they form.

Figure 6.20 compares the relative sizes of the atoms and ions for threemetals in Group 1A. For each of these elements, the ion is much smallerthan the atom. For example, the radius of a sodium ion (95 pm) is abouthalf the radius of a sodium atom (191 pm). When a sodium atom loses anelectron, the attraction between the remaining electrons and the nucleusis increased. The electrons are drawn closer to the nucleus. Also, metalsthat are representative elements tend to lose all their outermost electronsduring ionization. So the ion has one fewer occupied energy level.

The trend is the opposite for nonmetals like the halogens in Group 7A.For each of these elements, the ion is much larger than the atom. For exam-ple, the radius of a fluoride ion (133 pm) is more than twice the radius of afluorine atom (62 pm). As the number of electrons increases, the attractionof the nucleus for any one electron decreases.

Look back at Figure 6.19. From left to right across a period, two trendsare visible—a gradual decrease in the size of the positive ions followed by agradual decrease in the size of the negative ions. Figure 6.21 summarizesthe group and periodic trends in ionic size.

Figure 6.20 This diagram compares the relative sizes of atoms and ions for selected alkali metals and halogens. The data are given in picometers.Comparing and Contrasting What happens to the radius when an atom forms a cation? When an atom forms an anion?

Figure 6.21 The ionic radii for cations and anions decrease from left to right across periods and increase from top to bottom within groups.

LiLi

e

Group 1A Group 7A

e

e e

60 62 133

FF

Na Cl

ClNa

e e

181191

156

95

238 133

K

K

120

102

196

Br

Br

Siz

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incr

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Size of cations decreases

Trends in Ionic Size

Size of anions decreases


Trends in ElectronegativityDiscussLead a class discussion on periodic and group trends in electronegativities. Point out that electronegativity values help chemists predict the type of bond-ing that exists between atoms in com-pounds. Ask, Why aren't the noble gases included in a discussion on electronegativity? (because they form very few compounds) Which element is the most electronegative and which is the least electronegative? (fluorine; cesium) You may want to reference the electronegativity graphs in the Ele-ments Handbook on R7, R11, R15, R19, R25, R29, and R33.

FYIThe values for electronegativity are often based on values for ionization energy and electron affinity. Ionization energy is a measure of an atom’s ability to lose electrons. Electron affinity is a measure of an atom’s ability to gain electrons.

Download a worksheet on Electro–negativity for students to complete, and find additional teacher support from NSTA SciLinks.

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Answers to...Figure 6.20 The radius decreases for cations and increases for anions.

Checkpoint

They rarely form compounds.


Table 6.2

Electronegativity Values for Selected Elements

Trends in ElectronegativityIn Chapters 7 and 8, you will study two types of bonds that can exist incompounds. Electrons are involved in both types of bonds. There is a prop-erty that can be used to predict the type of bond that will form during areaction. This property is called electronegativity. Electronegativity is theability of an atom of an element to attract electrons when the atom is in acompound. Scientists use factors such as ionization energy to calculate val-ues for electronegativity.

Table 6.2 lists electronegativity values for representative elements inGroups 1A through 7A. The elements are arranged in the same order as in aperiodic table. The noble gases are omitted because they do not form manycompounds. The data in Table 6.2 is expressed in units called Paulings.Linus Pauling won a Nobel Prize in Chemistry for his work on chemicalbonds. He was the first to define electronegativity.

In general, electronegativity values decrease from top to bottomwithin a group. For representative elements, the values tend to increasefrom left to right across a period. Metals at the far left of the periodic tablehave low values. By contrast, nonmetals at the far right (excluding noblegases) have high values. The electronegativity values among the transitionmetals are not as regular.

The least electronegative element is cesium, with an electronegativityvalue of 0.7. It has the least tendency to attract electrons. When it reacts, ittends to lose electrons and form positive ions. The most electronegativeelement is fluorine, with a value of 4.0. Because fluorine has such a strongtendency to attract electrons, when it is bonded to any other element iteither attracts the shared electrons or forms a negative ion.

Checkpoint Why are values for noble gases omitted from Table 6.2?

For: Links on Electronegativity


H

2.1

Li Be B C N O F

1.0 1.5 2.0 2.5 3.0 3.5 4.0

Na Mg Al Si P S Cl

0.9 1.2 1.5 1.8 2.1 2.5 3.0

K Ca Ga Ge As Se Br

0.8 1.0 1.6 1.8 2.0 2.4 2.8

Rb Sr In Sn Sb Te I

0.8 1.0 1.7 1.8 1.9 2.1 2.5

Cs Ba Tl Pb Bi

0.7 0.9 1.8 1.9 1.9

178 Chapter 6


Summary of TrendsUse VisualsFigure 6.22 Point out that this diagram incorporates information from several diagrams earlier in the chapter. Ask stu-dents whether they find this or earlier diagrams more helpful, and why.

ASSESSEvaluate UnderstandingHave students compare two elements in the same group in terms of atomic radius, ionic radius, ionization energy, and electronegativity. Repeat the exer-cise with a metal and nonmetal from the same period. Have students write general statements to summarize the trends revealed by these comparisons.

ReteachReview the terms used in Figure 6.22. Then, use the periodic table and the terms to play a version of “I’m Thinkingof . . . .” For example, choose fluorine and say you are thinking of an element that has a very small atomic size and a very high electronegativity. Let students guess, then discuss the correct answer. Have students continue the game in small groups.

From top to bottom within a group, the effective nuclear charge remains nearly constant and the principal quantum number increases. From left to right in a period, the principal quantum number remains constant, but the effective nuclear charge increases.

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Section 6.3 Assessment16. Atomic size generally increases within a

group and decreases from left to right across a period.

17. Ions form when electrons are transferred between atoms.

18. decreases within a group and increases from left to right across a period

19. Anions are larger and cations are smaller than the atoms from which they form.

20. decreases from top to bottom within a group and increases from left to right across a period

21. The trends can be explained by variations in atomic structure.

22. sodium, aluminum, sulfur, chlorine; periodic trend

23. a. sodium b. phosphorus

178 Chapter 6

withChemASAP

Summary of TrendsFigure 6.22 shows the trends for atomic size, ionization energy, ionic size,and electronegativity in Groups 1A through 8A. These properties vary withingroups and across periods. The trends that exist among these proper-ties can be explained by variations in atomic structure. The increasein nuclear charge within groups and across periods explains many trends.Within groups an increase in shielding has a significant effect.


16. Key Concept How does atomic size change within groups and across periods?

17. Key Concept When do ions form?

18. Key Concept What happens to first ionization energy within groups and across periods?

19. Key Concept Compare the size of ions to the size of the atoms from which they form.

20. Key Concept How does electronegativity vary within groups and across periods?

21. Key Concept In general, how can the periodic trends displayed by elements be explained?

22. Arrange these elements in order of decreasing atomic size: sulfur, chlorine, aluminum, and sodium. Does your arrangement demonstrate a periodic trend or a group trend?

23. Which element in each pair has the larger first ionization energy?a. sodium, potassiumb. magnesium, phosphorus

Explaining Trends in Atomic Size Explain why the size of an atom tends to increase from top to bottom within a group. Explain why the size of an atom tends to decrease from left to right across a period.


Figure 6.22 Properties that vary within groups and across periods include atomic size, ionic size, ionization energy, electronegativity, nuclear charge, and shielding effect.Interpreting Diagrams Which properties tend to decrease across a period?

Nuclear charge increases

Shielding is constant

Atomic size decreases

Ionization energy increases

Electronegativity increases

Nu

clea

r ch

arg

e in

crea

ses

Sh

ield

ing

incr

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s

Ion

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ize

incr

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Ele

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Size of cations decreases Size of anions decreases

1A

2A 3A 4A 5A 6A 7A

8A


Small-ScaleLAB

Small-ScaleLAB

Periodicity in Three Dimensions

ObjectiveAfter completing this activity, students should be able to:• build concrete models to reinforce

periodic trends.• apply a procedure to a new variable

and design a model on their own.

Prep Time 10 minutesClass Time 40 minutes

Materials 96-well spot plates, straws, scissors, metric rulers, permanent fine-line markers

Advance Prep Straws with a 1/4-inch diameter fit snugly in the wells.

Teaching Tips• Students can use colored straws to

color code groups or periods.• If you don’t have spot plates, press a

lump of clay the size of a golf ball flat on a table with a block of wood. Stu-dents mark out a 1-cm square grid and insert the straws in the clay.

Expected Outcome Students pro-duce 3-D models for periodic trends.

Analyze and Conclude1. fluorine2. Electronegativity generally increases

from left to right along a period.3. Metals, which are on the left side of

the table, have lower electronegativ-ity values than nonmetals, which are on the right.

4. Electronegativity generally increases from bottom to top within a group. Except for boron, the rest of Group 3A shows a reverse in this trend.

5. Hydrogen is placed in Group 1A based on its electron configuration, but is classified as a nonmetal.

For Enrichment Have students use the data on page R41 to make a 3-D model of trends in atomic size for transition metals.

L2

L3You’re the Chemist1. Students divide the values of first ionization

energies by 300 and measure the appropri-ate length of straws.

2. Students must determine their own scale before they begin. Students often use two wells to represent both ionic and atomic radii. Other students cut a straw to a length that represents the larger radius of an atom and mark the straw to show the smaller radius of the corresponding cation.

3. The value for xenon is similar to iodine, which is consistent with the general trend. Based on this value, xenon appears to have the ability to attract electrons and form compounds.

Answers to. . .Figure 6.22 size of atoms and ions

Small-Scale Lab 179

Small-ScaleLAB

Small-ScaleLAB

Periodicity in Three Dimensions

PurposeTo build three-dimensional models for periodic trends.

Materials

• 96-well spot plate

• straws

• scissors

• metric ruler

• permanent fine-line marker

Procedure 1. Measure the depth of a well in the spot plate by

inserting a straw into a well and holding the straw upright as shown in the photograph. Make a mark on the straw at the point where the straw meets the surface of the plate. Measure the distance from the end of the straw to the mark in centimeters. Record this distance as well depth.

2. Cut the straw to a length that is 4.0 cm plus well depth. The straw will extend exactly 4.0 cm above the surface of the plate.

3. Fluorine has an electronegativity value of 4.0. On a scale of one cm equals one unit of electronegativity, the portion of the straw that extends above the surface of the plate represents the electronegativity value for fluorine. Using the same scale, cut straws to represent the electronegativity values for all the elements listed in Table 6.2. Remember to add the well depth to the electronegativity value before cutting a straw. As you cut the straws, mark each straw with the chemical symbol of the element that the straw represents.

4. Arrange the straws in the spot plate in rows and columns to match the locations of the elements in the periodic table.

5. Make a rough sketch of your completed model.

Analyze and ConcludeObserve your model and record the answers to the following questions below your sketch.

1. Which element represented in your model is the most electronegative?

2. Based on your model, what is the general trend in electronegativity from left to right across a period?

3. Relate the trend in electronegativity across a period to the location of metals and nonmetals in the periodic table.

4. What is the general trend in electronegativity within a group? Are there any notable exceptions?

5. Why do you think that the electronegativity value for hydrogen is so high given its location in the table?

You’re the ChemistThe following small-scale activities allow you to develop your own procedures and analyze the results.

1. Design It! Construct a similar 3-D model for first ionization energies. Use the data in Table 6.1. Use a scale of one cm equals 300 kJ/mol.

2. Design It! Design and construct a 3-D model that shows trends in atomic and ionic radii for the elements in Groups 1A and 7A. Devise a way to display both ionic and atomic radii in the same model.

3. Analyze it! Xenon has an electronegativity value of 2.6. Cut and place a straw in your first model to represent xenon. Does xenon support the trend for electronegativity across a period? Is xenon likely to form compounds? Explain your answers.

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