CHAPTER 5 The Periodic Law · mass, is the basis for the organization of the periodic table....

C H A P T E R 5

The Periodic Law

The physical and chemical properties of the elements are periodic functions of their atomic numbers.

Copyright © by Holt, Rinehart and Winston. All rights reserved.


History of thePeriodic Table

I magine the confusion among chemists during the middle of thenineteenth century. By 1860, more than 60 elements had been discov-ered. Chemists had to learn the properties of these elements as well asthose of the many compounds that they formed—a difficult task. Andto make matters worse, there was no method for accurately determin-ing an element’s atomic mass or the number of atoms of an element ina particular chemical compound. Different chemists used differentatomic masses for the same elements, resulting in different composi-tions being proposed for the same compounds. This made it nearlyimpossible for one chemist to understand the results of another.

In September 1860, a group of chemists assembled at the FirstInternational Congress of Chemists in Karlsruhe, Germany, to settle theissue of atomic mass as well as some other matters that were makingcommunication difficult. At the Congress, Italian chemist StanislaoCannizzaro presented a convincing method for accurately measuringthe relative masses of atoms. Cannizzaro’s method enabled chemists toagree on standard values for atomic mass and initiated a search for rela-tionships between atomic mass and other properties of the elements.

Mendeleev and Chemical Periodicity

When the Russian chemist Dmitri Mendeleev heard about the newatomic masses discussed at Karlsruhe, he decided to include the newvalues in a chemistry textbook that he was writing. In the book,Mendeleev hoped to organize the elements according to their proper-ties. He went about this much as you might organize information for aresearch paper. He placed the name of each known element on a card,together with the atomic mass of the element and a list of its observedphysical and chemical properties. He then arranged the cards accordingto various properties and looked for trends or patterns.

Mendeleev noticed that when the elements were arranged in orderof increasing atomic mass, certain similarities in their chemical proper-ties appeared at regular intervals. Such a repeating pattern is referred toas periodic. The second hand of a watch, for example, passes over anygiven mark at periodic, 60-second intervals. The circular waves createdby a drop of water hitting a water surface are also periodic.

T H E P E R I O D I C L A W 123

SECTION 5-1

OBJECTIVES

Explain the roles of Mendeleevand Moseley in the develop-ment of the periodic table.

Describe the modern periodictable.

Explain how the periodic law can be used to predict the physical and chemicalproperties of elements.

Describe how the elementsbelonging to a group of theperiodic table are interrelatedin terms of atomic number.

FIGURE 5-1 The regularly spaced water waves represent a simple periodic pattern.

NSTA

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Mendeleev created a table in which elements with similar propertieswere grouped together—a periodic table of the elements. His first peri-odic table, shown in Figure 5-2, was published in 1869. Note thatMendeleev placed iodine, I (atomic mass 127), after tellurium,Te (atom-ic mass 128). Although this contradicted the pattern of listing the ele-ments in order of increasing atomic mass, it allowed Mendeleev to placetellurium in a group of elements with which it shares similar properties.Reading horizontally across Mendeleev’s table, this group includes oxy-gen, O, sulfur, S, and selenium, Se. Iodine could also, then, be placed inthe group it resembles chemically, which includes fluorine, F, chlorine,Cl, and bromine, Br.

Mendeleev’s procedure left several empty spaces in his periodictable (see Figure 5-2). In 1871, the Russian chemist boldly predicted theexistence and properties of the elements that would fill three of thespaces. By 1886, all three elements had been discovered. Today theseelements are known as scandium, Sc, gallium, Ga, and germanium, Ge.Their properties are strikingly similar to those predicted by Mendeleev.

C H A P T E R 5124

FIGURE 5-2 In his first pub-lished periodic table, Mendeleevarranged the elements in verticalperiods according to relativeatomic mass. The atomic massfor each element is indicated by the number following the element’s symbol. The unknownelements indicated by questionmarks at estimated atomic mass-es 45, 68, and 70 were later iden-tified as scandium, Sc, gallium,Ga, and germanium, Ge.


The success of Mendeleev’s predictions persuaded most chemists toaccept his periodic table and earned him credit as the discoverer of theperiodic law. Two questions remained, however. (1) Why could most ofthe elements be arranged in the order of increasing atomic mass but afew could not? (2) What was the reason for chemical periodicity?

Moseley and the Periodic Law

The first question was not answered until more than 40 years afterMendeleev’s first periodic table was published. In 1911, the English sci-entist Henry Moseley, who was working with Ernest Rutherford, exam-ined the spectra of 38 different metals. When analyzing his data,Moseley discovered a previously unrecognized pattern. The elements inthe periodic table fit into patterns better when they were arranged inincreasing order according to nuclear charge, or the number of protonsin the nucleus. Moseley’s work led to both the modern definition ofatomic number and the recognition that atomic number, not atomicmass, is the basis for the organization of the periodic table.

Moseley’s discovery was consistent with Mendeleev’s ordering of theperiodic table by properties rather than strictly by atomic mass. Forexample, according to Moseley, tellurium, with an atomic number of 52,belongs before iodine, which has an atomic number of 53. Today,Mendeleev’s principle of chemical periodicity is correctly stated in whatis known as the periodic law: The physical and chemical properties of theelements are periodic functions of their atomic numbers. In other words,when the elements are arranged in order of increasing atomic number,elements with similar properties appear at regular intervals.

The Modern Periodic Table

The periodic table has undergone extensive change since Mendeleev’stime (see Figure 5-6 on pages 130–131). Chemists have discovered newelements and, in more recent years, synthesized new ones in the lab-oratory. Each of the more than 40 new elements, however, can beplaced in a group of other elements with similar properties. Theperiodic table is an arrangement of the elements in order of theiratomic numbers so that elements with similar properties fall in the samecolumn, or group.

The Noble GasesPerhaps the most significant addition to the periodic table came withthe discovery of the noble gases. In 1894, English physicist John WilliamStrutt (Lord Rayleigh) and Scottish chemist Sir William Ramsay dis-covered argon, Ar, a gas in the atmosphere that had previously escapednotice because of its total lack of chemical reactivity. Back in 1868,


FIGURE 5-3 The noble gases, alsoknown as the Group 18 elements,are all rather unreactive. As you willread, the reason for this low reactiv-ity also accounts for the special placeoccupied by the noble gases in theperiodic table.

7

N

15

P

8

O9

F10

Ne

2

He

16

S17

Cl18

Ar

33

As34

Se35

Br36

Kr

51

Sb52

Te53

I54

Xe

83

Bi84

Po85

At86

Rn

NSTA

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Differencein atomicnumbers

Gro

up

18

Gro

up

1

He

Ne

Ar

Kr

Xe

Rn

Li

Na

K

Rb

Cs

Fr

Elementand atomicnumber

2

10

18

36

54

86

8

8

18

18

32

8

8

18

18

32

3

11

19

37

55

87

another noble gas, helium, He, had been discovered as a component ofthe sun, based on the emission spectrum of sunlight. In 1895, Ramsayshowed that helium also exists on Earth.

In order to fit argon and helium into the periodic table, Ramsay pro-posed a new group. He placed this group between the groups nowknown as Group 17 (the fluorine family) and Group 1 (the lithium fam-ily). In 1898, Ramsay discovered two more noble gases to place in hisnew group, krypton, Kr, and xenon, Xe. The final noble gas, radon, Rn,was discovered in 1900 by the German scientist Friedrich Ernst Dorn.

The LanthanidesThe next step in the development of the periodic table was completedin the early 1900s. It was then that the puzzling chemistry of the lan-thanides was finally understood. The lanthanides are the 14 elementswith atomic numbers from 58 (cerium, Ce) to 71 (lutetium, Lu). Becausethese elements are so similar in chemical and physical properties, theprocess of separating and identifying them was a tedious task thatrequired the effort of many chemists.

The ActinidesAnother major step in the development of the periodic table was thediscovery of the actinides. The actinides are the 14 elements with atomicnumbers from 90 (thorium, Th) to 103 (lawrencium, Lr). The lan-thanides and actinides belong in Periods 6 and 7, respectively, of theperiodic table, between the elements of Groups 3 and 4. To save space,the lanthanides and actinides are usually set off below the main portionof the periodic table, as shown in Figure 5-6 on pages 130–131.

PeriodicityPeriodicity with respect to atomic number can be observed in any groupof elements in the periodic table. Consider the noble gases of Group 18.The first noble gas is helium, He. It has an atomic number of 2. The ele-ments following helium in atomic number have completely differentproperties until the next noble gas, neon, Ne, which has an atomic num-ber of 10, is reached. The remaining noble gases in order of increasingatomic number are argon (Ar, atomic number 18), krypton (Kr, atomicnumber 36), xenon (Xe, atomic number 54), and radon (Rn, atomicnumber 86). The differences in atomic number between successivenoble gases are shown in Figure 5-4. Also shown in Figure 5-4 areatomic-number differences between the elements of Group 1, which are all solid, silvery metals. As you can see, the differences in atomicnumber between the Group 1 metals follow the same pattern as the differences in atomic number between the noble gases.

Starting with the first member of Groups 13–17, a similar periodicpattern is repeated. The atomic number of each successive element is 8,18, 18, and 32 higher than the atomic number of the element above it.In Section 5-2, you will see that the second mystery presented byMendeleev’s periodic table—the reason for periodicity—is explainedby the arrangement of the electrons around the nucleus.

C H A P T E R 5126

FIGURE 5-4 In each of Groups 1and 18, the differences between theatomic numbers of successive ele-ments are 8, 8, 18, 18, and 32, respec-tively. Groups 2 and 13–17 follow asimilar pattern.



SECTION REVIEW

1. a. Who is credited with developing a method thatled to the determination of standard relativeatomic masses?

b. Who discovered the periodic law?c. Who established atomic numbers as the basis

for organizing the periodic table?

2. State the periodic law.

3. Name three sets of elements added to the periodictable after Mendeleev’s time.

4. How do the atomic numbers of the elements with-in each of Groups 1, 2, and 13–18 of the periodictable vary? (Refer to Figure 5-4 as a guide.)

• index cards

Question

Materials

DiscussionProcedure

Designing Your Own Periodic Table

Can you design your own periodictable using information similar tothat available to Mendeleev?

as you think Mendeleev mighthave done.

1. Keeping in mind that the information you have is similarto that available to Mendeleev in 1869, answer the followingquestions.

a. Why are atomic masses giveninstead of atomic numbers?

b. Can you identify each ele-ment by name?

2. How many groups of elements,or families, are in your periodictable? How many periods, orseries, are in the table?

3. Predict the characteristics of anymissing elements.When you havefinished, check your work usingyour separate list of elements anda periodic table.

1. Write down the information avail-able for each element on separateindex cards. The following infor-mation is appropriate: a letter ofthe alphabet (A, B, C, etc.) to

identify each element; atomicmass; state; density; melting

point; boiling point; andany other readily observ-

able physical proper-ties. Do not write the

name of the ele-ment on the index

card, but keep a sepa-rate list indicating the letters

you have assigned to eachelement.

2. Organize the cards for the elements in a logical pattern


ElectronConfiguration and the Periodic Table

T he Group 18 elements of the periodic table (the noble gases)undergo few chemical reactions. This stability results from the gases’special electron configurations. Helium’s highest occupied level, the 1sorbital, is completely filled with electrons. And the highest occupiedlevels of the other noble gases contain stable octets. Generally the elec-tron configuration of an atom’s highest occupied energy level governsthe atom’s chemical properties.

Periods and Blocks of the Periodic Table

While the elements are arranged vertically in the periodic table ingroups that share similar chemical properties, they are also organizedhorizontally in rows, or periods. (As shown in Figure 5-6, there are a totalof seven periods of elements in the modern periodic table.) As can beseen in Table 5-1, the length of each period is determined by the numberof electrons that can occupy the sublevels being filled in that period.

C H A P T E R 5128

SECTION 5-2

OBJECTIVES

Describe the relationshipbetween electrons in sub-levels and the length of eachperiod of the periodic table.

Locate and name the fourblocks of the periodic table.Explain the reasons for these names.

Discuss the relationshipbetween group configurationsand group numbers.

Describe the locations in the periodic table and thegeneral properties of the alkali metals, the alkaline-earth metals, the halogens,and the noble gases.

Period Number of Sublevels innumber elements in period order of filling

1 2 1s

2 8 2s 2p

3 8 3s 3p

4 18 4s 3d 4p

5 18 5s 4d 5p

6 32 6s 4f 5d 6p

7 32 7s 5f 6d, etc.

TABLE 5-1 Relationship Between Period Length andSublevels Being Filled in the Periodic Table


In the first period, the 1s sublevel is being filled. The 1s sublevel canhold a total of two electrons. Therefore, the first period consists of twoelements—hydrogen and helium. In the second period, the 2s sub-level, which can hold two electrons, and the 2p sublevel, which canhold six electrons, are being filled. Consequently, the second periodtotals eight elements. Similarly, filling of the 3s and 3p sublevelsaccounts for the eight elements of the third period. Filling 3d and 4dsublevels in addition to the s and p sublevels adds 10 elements to boththe fourth and fifth periods. Therefore, each of these periods totals 18 elements. Filling 4f sublevels in addition to s, p, and d sublevelsadds 14 elements to the sixth period, which totals 32 elements. And asnew elements are created, the 29 known elements in Period 7 could, intheory, be extended to 32.

The period of an element can be determined from the element’selectron configuration. For example, arsenic, As, has the electronconfiguration [Ar]3d 104s24p3. The 4 in 4p3 indicates that arsenic’shighest occupied energy level is the fourth energy level. Arsenic istherefore in the fourth period in the periodic table. The period andelectron configuration for each element can be found in the periodictable on pages 130–131.

Based on the electron configurations of the elements, the periodictable can be divided into four blocks, the s, p, d, and f blocks. This divi-sion is illustrated in Figure 5-5. The name of each block is determinedby whether an s, p, d, or f sublevel is being filled in successive elementsof that block.


FIGURE 5-5 Based on the elec-tron configurations of the elements,the periodic table can be subdividedinto four sublevel blocks.

1

H

3

Li11

Na19

K37

Rb55

Cs87

Fr

4

Be12

Mg20

Ca38

Sr56

Ba88

Ra

5

B13

Al31

Ga49

In81

Tl113

6

C14

Si32

Ge50

Sn82

Pb114

Uuq

7

N15

P33

As51

Sb83

Bi115

8

O16

S34

Se52

Te84

Po116

Uuh

9

F17

Cl35

Br53

I85

At117

21

Sc39

Y57

La89

Ac

22

Ti40

Zr72

Hf104

Rf

23

V41

Nb73

Ta105

Db

24

Cr42

Mo74

W106

Sg

25

Mn43

Tc75

Re107

Bh

26

Fe44

Ru76

Os108

Hs

27

Co45

Rh77

Ir109

Mt

28

Ni46

Pd78

Pt110

Uun

29

Cu47

Ag79

Au111

Uuu

30

Zn48

Cd80

Hg112

Uub

Group 1

2

He10

Ne18

Ar36

Kr54

Xe86

Rn118

Uuo

66

Dy98

Cf

67

Ho99

Es

68

Er100

Fm

69

Tm101

Md

70

Yb102

No

58

Ce90

Th

59

Pr91

Pa

60

Nd92

U

61

Pm93

Np

62

Sm94

Pu

63

Eu95

Am

64

Gd96

Cm

65

Tb97

Bk

71

Lu103

Lr

Group 18

Group 2

Group 3 Group 4 Group 5 Group 6 Group 7 Group 8 Group 9 Group 10 Group 11 Group 12

Group 13 Group 14 Group 15 Group 16 Group 17

s-block elements

p-block elements

d-block elements

f-block elements

Sublevel Blocks of the Periodic Table


1

HHydrogen

1.007941s1

3Li

Lithium6.941

[He]2s1

11Na

Sodium22.989768

[Ne]3s1

19K

Potassium39.0983[Ar]4s1

37Rb

Rubidium85.4678[Kr]5s1

55Cs

Cesium132.90543

[Xe]6s1

87Fr

Francium(223.0197)

[Rn]7s1

58Ce

Cerium140.115

[Xe]4f15d16s2

90Th

Thorium232.0381[Rn]6d27s2

59Pr

Praseodymium140.908

[Xe]4f 36s2

91Pa

Protactinium231.03588

[Rn]5f 26d17s2

60Nd

Neodymium144.24

[Xe]4f 46s2

92U

Uranium238.0289

[Rn]5f 36d17s2

61Pm

Promethium(144.9127)[Xe]4f 56s2

93Np

Neptunium(237.0482)

[Rn]5f 46d17s2

62Sm

Samarium150.36

[Xe]4f 66s2

94Pu

Plutonium244.0642[Rn]5f 67s2

4Be

Beryllium9.012182

[He]2s2

12Mg

Magnesium24.3050[Ne]3s2

20Ca

Calcium40.078[Ar]4s2

38Sr

Strontium87.62[Kr]5s2

56Ba

Barium137.327[Xe]6s2

88Ra

Radium(226.0254)

[Rn]7s2

21Sc

Scandium44.955910[Ar]3d14s2

39Y

Yttrium88.90585[Kr]4d15s2

57La

Lanthanum138.9055[Xe]5d16s2

89Ac

Actinium(227.0278)[Rn]6d17s2

22Ti

Titanium47.88

[Ar]3d24s2

40Zr

Zirconium91.224

[Kr]4d25s2

72Hf

Hafnium178.49

[Xe]4f145d26s2

104Rf

Rutherfordium(261.11)

[Rn]5f146d27s2

23V

Vanadium50.9415

[Ar]3d34s2

41Nb

Niobium92.90638[Kr]4d45s1

73Ta

Tantalum180.9479

[Xe]4f145d36s2

105Db

Dubnium(262.114)

[Rn]5f146d37s2

24Cr

Chromium51.9961

[Ar]3d54s1

42Mo

Molybdenum95.94

[Kr]4d55s1

74W

Tungsten183.84

[Xe]4f145d46s2

106Sg

Seaborgium(263.118)

[Rn]5f146d47s2

25Mn

Manganese54.93805[Ar]3d54s2

43Tc

Technetium(97.9072)[Kr]4d65s1

75Re

Rhenium186.207

[Xe]4f145d56s2

107Bh

Bohrium(262.12)

[Rn]5f146d57s2

26FeIron

55.847[Ar]3d64s2

44Ru

Ruthenium101.07

[Kr]4d75s1

76Os

Osmium190.23

[Xe]4f145d66s2

108Hs

Hassium(265)†

[Rn]5f146d67s2

27Co

Cobalt58.93320[Ar]3d74s2

45Rh

Rhodium102.906

[Kr]4d85s1

77Ir

Iridium192.22

[Xe]4f145d76s2

109Mt

Meitnerium(266)†

[Rn]5f146d77s2

Group 1 Group 2

Group 3 Group 4 Group 5 Group 6 Group 7 Group 8 Group 9

2

3

4

5

6

7

Perio

d

* Temporary name not yet approved byIUPAC.

† Estimated from currently available IUPAC data.

1 Atomic number

Symbol

Name

Average atomic mass

Electron configuration

6C

Carbon12.011

[He]2s22p2

Key:

C H A P T E R 5130

Periodic Table of the Elements

FIGURE 5-6 In the common periodic table, the elements are arranged in vertical groups and in horizontal periods.

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[Xe]4f 145d106s26p6

[Kr]4d 105s25p6

[Ar]3d 104s24p6

[Ne]3s 23p6

1s 2

[He]2s 22p5[He]2s 22p4[He]2s 22p3[He]2s 22p2[He]2s 22p1

[Ne]3s 23p5[Ne]3s 23p4[Ne]3s 23p3[Ne]3s 23p2[Ne]3s 23p1

[Ar]3d 104s24p5[Ar]3d 104s24p4[Ar]3d 104s24p3[Ar]3d 104s24p2[Ar]3d 104s24p1[Ar]3d 104s2[Ar]3d 104s1[Ar]3d 84s2

[Kr]4d 105s25p5[Kr]4d 105s25p4[Kr]4d 105s25p3[Kr]4d 105s25p2[Kr]4d 105s25p1[Kr]4d 105s2[Kr]4d 105s1[Kr]4d 105s0

[Xe]4f 145d106s26p5[Xe]4f 145d106s26p4[Xe]4f 145d106s26p3[Xe]4f 145d106s26p2[Xe]4f 145d106s26p1[Xe]4f 145d106s2

28Ni

Nickel58.6934

29Cu

Copper63.546

30ZnZinc65.39

31Ga

Gallium69.723

32Ge

Germanium72.61

33As

Arsenic74.92159

34Se

Selenium78.96

35Br

Bromine79.904

36Kr

Krypton83.80

13Al

Aluminum26.981539

14Si

Silicon28.0855

15P

Phosphorus30.9738

16S

Sulfur32.066

17Cl

Chlorine35.4527

18Ar

Argon39.948

5B

Boron10.811

6C

Carbon12.011

7N

Nitrogen14.00674

8O

Oxygen15.9994

9F

Fluorine18.9984032

10NeNeon20.1797

[He]2s 22p6

2He

Helium4.002602

46Pd

Palladium106.42

47AgSilver

107.8682

48Cd

Cadmium112.411

49In

Indium114.818

50SnTin

118.710

51Sb

Antimony121.757

52Te

Tellurium127.60

53I

Iodine126.904

54Xe

Xenon131.29

[Xe]4f 145d96s1

78Pt

Platinum195.08

[Rn]5f 146d 97s1

110Uun*

Ununnilium(269)†

[Rn]5f 146d 107s1

111Uuu*

Unununium(272)†

[Rn]5f 146d 107s2

112Uub*

Ununbium(277)†

113

[Rn]5f 146d 107s27p2

114Uuq*

Ununquadium(285)†

115

[Rn]5f 146d 107s27p4

116Uuh*

Ununhexium(289)†

117

[Rn]5f 146d 107s27p6

118Uuo*

Ununoctium(293)†

[Xe]4f 145d106s1

79AuGold

196.96654

80Hg

Mercury200.59

81Tl

Thallium204.3833

82PbLead207.2

83Bi

Bismuth208.98037

84Po

Polonium(208.9824)

85At

Astatine(209.9871)

86Rn

Radon(222.0176)


Group 18

Group 10 Group 11 Group 12

Metals

Alkali metals

Alkaline earth metals

Transition metals

Other metals

Metalloids

Semiconductors

Nonmetals

Halogens

Other nonmetals

Noble gases

1

2

3

4

5

6

7

Period

63Eu

Europium151.966

95Am

Americium(243.0614)

64Gd

Gadolinium157.25

96CmCurium

(247.0703)

65Tb

Terbium158.92534

97Bk

Berkelium(247.0703)

66Dy

Dysprosium162.50

98Cf

Californium(251.0796)

67

Holmium164.930

99

Einsteinium(252.083)

68

Erbium167.26

100

Fermium(257.0951)

69

Thulium168.93421

101

Mendelevium(258.10)

70

Ytterbium173.04

102

Nobelium(259.1009)

71

Lutetium174.967

[Xe]4f 145d16s2[Xe]4f 146s2[Xe]4f 136s2[Xe]4f 126s2[Xe]4f 116s2[Xe]4f 106s2[Xe]4f 96s2[Xe]4f 75d16s2[Xe]4f 76s2

103

Lawrencium(262.11)

[Rn]5f 146d17s2[Rn]5f 147s2[Rn]5f 137s2[Rn]5f 127s2[Rn]5f 117s2[Rn]5f 107s2[Rn]5f 97s2[Rn]5f 76d17s2[Rn]5f 77s2

Ho

Es

Er

Fm

Tm

Md

Yb

No

Lu

Lr


The atomic masses listed in this table reflect the precision of current measurements. (Values listed inparentheses are those of the element’s most stable or most common isotope.) In calculations throughout the text, however, atomic masses have been rounded to two places to the right of the decimal.


The s-Block Elements: Groups 1 and 2The elements of the s block are chemically reactive metals. The Group 1metals are more reactive than those of Group 2. The outermost energylevel in an atom of each Group 1 element contains a single s electron.For example, the configurations of lithium and sodium are [He]2s1 and[Ne]3s1, respectively. As you will learn in Section 5-3, the ease withwhich the single electron is lost helps to make the Group 1 metalsextremely reactive. Using n for the number of the highest occupiedenergy level, the outer, or group, configurations of the Group 1 and 2elements are written ns1 and ns2, respectively. For example, the config-uration of Na is [Ne]3s1, so the group configuration is written ns1,where n � 3.

The elements of Group 1 of the periodic table (lithium, sodium, potas-sium, rubidium, cesium, and francium) are known as the alkali metals. Intheir pure state, all of the alkali metals have a silvery appearance andare soft enough to cut with a knife. However, because they are so reac-tive, alkali metals are not found in nature as free elements. They com-bine vigorously with most nonmetals. And they react strongly withwater to produce hydrogen gas and aqueous solutions of substancesknown as alkalis. Because of their extreme reactivity with air or mois-ture, alkali metals are usually stored in kerosene. Proceeding down thecolumn, the elements of Group 1 melt at successively lower tempera-tures.

The elements of Group 2 of the periodic table (beryllium, magnesium,calcium, strontium, barium, and radium) are called the alkaline-earthmetals. Atoms of alkaline-earth metals contain a pair of electrons intheir outermost s sublevel. Consequently, the group configuration forGroup 2 is ns2. The Group 2 metals are harder, denser, and strongerthan the alkali metals. They also have higher melting points. Althoughthey are less reactive than the alkali metals, the alkaline-earth metalsare also too reactive to be found in nature as free elements.

Hydrogen and HeliumBefore discussing the other blocks of the periodic table, let’s considertwo special cases in the classification of the elements—hydrogen andhelium. Hydrogen has an electron configuration of 1s1, but despite thens1 configuration, it does not share the same properties as the elements

C H A P T E R 5132

FIGURE 5-8 Calcium, an alkaline-earth metal, is too reactive to befound in nature in its pure state (a).Instead, it exists in compounds, suchas in the minerals that make up marble (b).

FIGURE 5-7 (a) Like other alkalimetals, potassium reacts so stronglywith water that (b) it must be storedin kerosene or oil to prevent it fromreacting with moisture in the air.

(a)

(b)

(a) (b)


of Group 1. Although it is located above the Group 1 elements in manyperiodic tables, hydrogen is a unique element, with properties that donot closely resemble those of any group.

Like the Group 2 elements, helium has an ns2 group configuration.Yetit is part of Group 18. Because its highest occupied energy level is filledby two electrons, helium possesses special chemical stability, exhibitingthe unreactive nature of a Group 18 element. By contrast, the Group 2metals have no special stability; their highest occupied energy levels arenot filled because each metal has an empty available p sublevel.


a. Without looking at the periodic table, give the group, period, and block in which the element with the electron configuration [Xe]6s2 is located.

b. Without looking at the periodic table, write the electron configuration for the Group 1element in the third period. Is this element likely to be more active or less active thanthe element described in (a)?

a. The element is in Group 2, as indicated by the group configuration of ns2. It is in thesixth period, as indicated by the highest principal quantum number in its configuration,6. The element is in the s block.

b. In a third-period element, the highest occupied energy level is the third main energylevel, n = 3. The 1s, 2s, and 2p sublevels are completely filled (see Table 5-1). A Group 1element has a group configuration of ns1, which indicates a single electron in its highest s sublevel. Therefore, this element has the following configuration.

1s22s22p63s1 or [Ne]3s1

Because it is in Group 1 (the alkali metals), this element is likely to be more reactivethan the element described in (a), which is in Group 2 (the alkaline-earth metals).

SOLUTION

SAMPLE PROBLEM 5- 1

Answer1. Group 1, fifth period, s block

2. a. ns2

b. 1s22s22p63s23p64s2

c. Ca, [Ar]4s2

d. The element is in Group 2,so it is probably less reactivethan the Group 1 element ofthe same period.

1. Without looking at the periodic table, give thegroup, period, and block in which the elementwith the electron configuration [Kr]5s1 is located.

2. a. Without looking at the periodic table, write thegroup configuration for the Group 2 elements.

b. Without looking at the periodic table, writethe complete electron configuration for theGroup 2 element in the fourth period.

c. Refer to Figure 5-6 to identify the elementdescribed in (b). Then write the element’snoble-gas notation.

d. How does the reactivity of the element in (b)compare with the reactivity of the element inGroup 1 of the same period?

PRACTICE

NSTA

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TOPIC: Alkaline-earth metalsGO TO: www.scilinks.orgsci LINKS CODE: HC2054


C H A P T E R 5134

FIGURE 5-9 The diagram shows the electron configuration ofscandium, Sc, the Group 3 elementof the fourth period. In general, the (n − 1)d sublevel in Groups 3–12 isoccupied by electrons after the nssublevel is filled.

Ener

gy

1s2s 2p

3s 3p

4s 3d

FIGURE 5-10 Mercury, tungsten,and vanadium are transition elements.Locate them in the d block of theperiodic table on page 129.

Mercury Tungsten Vanadium

The d-Block Elements: Groups 3–12For energy level n, there are n possible sublevels, so the d sublevel firstappears when n � 3. This 3d sublevel is slightly higher in energy than the4s sublevel, so these are filled in the order 4s3d (see Figure 5-9). Thisorder of filling is also seen for higher values of n. Each d sublevel consistsof five orbitals with a maximum of two electrons each, or up to 10 elec-trons possible in each d sublevel. In addition to the two ns electrons ofGroup 2, atoms of the Group 3 elements each have one electron in the dsublevel of the (n − 1) energy level. The group configuration for Group 3is therefore (n − 1)d1ns2. Atoms of the Group 12 elements have 10 elec-trons in the d sublevel plus two electrons in the ns sublevel. The groupconfiguration for Group 12 is (n − 1)d10ns2.

As you read in Chapter 4, some deviations from orderly d sublevelfilling occur in Groups 4–11. As a result, elements in these d-blockgroups, unlike those in s-block and p-block groups, do not necessarilyhave identical outer electron configurations. For example, in Group 10,nickel, Ni, has the electron configuration [Ar]3d84s2. Palladium, Pd, hasthe configuration [Kr]4d105s0. And platinum, Pt, has the configuration[Xe]4f 145d96s1. Notice, however, that in each case the sum of the outer sand d electrons is equal to the group number.

The d-block elements are metals with typical metallic properties andare often referred to as transition elements. They are good conductors ofelectricity and have a high luster. They are typically less reactive thanthe alkali metals and the alkaline-earth metals. Some are so unreactivethat they do not easily form compounds, existing in nature as free ele-ments. Palladium, platinum, and gold are among the least reactive of allthe elements. Some d-block elements are shown in Figure 5-10.



C H E M I C A L C O M M E N T A R Y

W elcome to the PeriodicKingdom. This is a land of theimagination, but it is closer toreality than it appears to be. Thisis the kingdom of the chemicalelements, the substances fromwhich everything tangible is made.It is not an extensive country, forit consists of only a hundred or soregions . . . yet it accounts foreverything material in our actualworld. From the hundred elementsthat are at the center of our story,all planets, rocks, vegetation, andanimals are made. These elementsare the basis of the air, the oceans,and the Earth itself. We stand onthe elements, we eat the elements,we are the elements. Because ourbrains are made up of elements,even our opinions are, in a sense,properties of the elements andhence inhabitants of the kingdom.

Even from . . . far above thecountry, we can see broad fea-tures of the landscape. There arethe glittering, lustrous regionsmade up of metals and lyingtogether in what we shall call theWestern Desert. This desert isbroadly uniform, but there is asubtlety of shades, indicating avariety of characteristics. Hereand there are gentle splashes ofcolor, such as the familiar glint of gold and the blush of copper.How remarkable it is that thesedesert lands make up so much

of the kingdom . . . yet the kingdom supplies such luxuriance in the real world!

Generally speaking,the Western Desert wasexplored and exploited fromeast to west; that is, technologyand industry made use of them in that order. Copper displacedstone to give us the Bronze Age.Then, as the explorers pressedwestward, applying ever morevigorous means of discovery, theyencountered iron and used it tofabricate more effective weap-onry . . . The strongest statesenjoyed freedom from constantaggression, and this gave themtime for scholarship; thus in duecourse explorers were able topenetrate into more distant western regions of the desert.

Deep in the Western Desert,in an isthmuslike zone runningfrom zinc on the east to scandiumon the west, they finally stumbledupon titanium, a remarkable prizeindeed. Titanium has exactly theproperties that a society bent onhigh technology needs if it is to

take to the skies: this is a metalthat is tough and resistant to corrosion, yet light, and it is typical of its part of the WesternDesert. Titanium and its neigh-bors vanadium and molybdenum,in alliance with iron, form thedurable steels that enable us tochop through stone and build ona massive scale. It is a remarkablefeature of the Periodic Kingdomthat the Isthmus . . . has providedso many of the workhorses of our society, and that its membersso readily form alliances with one another.

Reading for MeaningWhat does Atkins mean when he says that metals of the Isthmusform “alliances” with one an-other? What does this implyabout these elements?

Read FurtherAtkins remarks that titanium hasproperties that are well-suited fora high-tech society. Find out howwe obtain titanium, and list fiveways that it is used in our society.

NorthernShore Hydrogen

OxygenNitrogen

Carbon

Helium

Sulfur

Bromine

No

ble

gas

es

Gold MercurySouthern Shore

Southern Island

Alk

ali m

etal

s

Alk

alin

e-ea

rth

met

als

Western Desert

N

S

W E

Lanthanides

Actinides

Hal

og

ens

Sc Ti V Cr Mn Fe Co Ni Cu Zn

Wes

tern

rec

tan

gle

East

ern

rec

tan

gle

The Wild KingdomFrom The Periodic Kingdom: A Journey Into the World of the Chemical Elements by P. W. Atkins

The general layout of the Periodic Kingdom.

The Isthmus


C H A P T E R 5136

An element has the electron configuration [Kr]4d55s1. Without looking at the periodictable, identify the period, block, and group in which this element is located. Then consultthe periodic table to identify this element and the others in its group.

The number of the highest occupied energy level is five, so the element is in the fifthperiod. There are five electrons in the d sublevel. This means that the d sublevel is incom-pletely filled because it can hold 10 electrons. Therefore, the element is in the d block. Ford-block elements, the number of electrons in the ns sublevel (1) plus the number of elec-trons in the (n − 1)d sublevel (5) equals the group number, 6. This is the Group 6 element in the fifth period. The element is molybdenum. The others in Group 6 are chromium,tungsten, and seaborgium.

SOLUTION

SAMPLE PROBLEM 5- 2

Answer1. fourth period, d block,

Group 10

2. a. 4d105s2

b. Cd, [Kr]4d105s2

1. Without looking at the periodic table, identifythe period, block, and group in which the ele-ment with the electron configuration [Ar]3d84s2

is located.

2. a. Without looking at the periodic table, writethe outer electron configuration for theGroup 12 element in the fifth period.

b. Refer to the periodic table to identify theelement described in (a) and to write the ele-ment’s noble-gas notation.

PRACTICE

The p-Block Elements: Groups 13–18The p-block elements consist of all the elements of Groups 13–18 excepthelium. Electrons add to a p sublevel only after the s sublevel in the sameenergy level is filled.Therefore, atoms of all p-block elements contain twoelectrons in the ns sublevel. The p-block elements together with the s-blockelements are called the main-group elements. For Group 13 elements, theadded electron enters the np sublevel, giving a group configuration ofns2np1. Atoms of Group 14 elements contain two electrons in the p sub-level, giving ns2np2 for the group configuration. This pattern continues inGroups 15–18. In Group 18, the stable noble-gas configuration of ns2np6

is reached. The relationships among group numbers and electron config-urations for all the groups are summarized in Table 5-2.

For atoms of p-block elements, the total number of electrons in thehighest occupied level is equal to the group number minus 10. For exam-ple, bromine is in Group 17. It has 17 − 10 = 7 electrons in its highestenergy level. Because atoms of p-block elements contain two electronsin the ns sublevel, we know that bromine has five electrons in its outerp sublevel. The electron configuration of bromine is [Ar]3d104s24p5.

The properties of elements of the p block vary greatly. At its right-hand end, the p block includes all of the nonmetals except hydrogen andhelium. All six of the metalloids (boron, silicon, germanium, arsenic,


FIGURE 5-11 Fluorine, chlorine,bromine, and iodine are members of Group 17 of the periodic table,also known as the halogens. Locatethe halogens in the p block of theperiodic table on page 129.


Group number Group configuration Block Comments

1, 2 ns1, 2 s One or two electrons in ns sublevel

3–12 (n − 1)d1–10ns0–2 d Sum of electrons in ns and (n − 1)dlevels equals group number

13–18 ns2np1–6 p Number of electrons in np sublevel equalsgroup number minus 12

TABLE 5-2 Relationships Among Group Numbers, Blocks, and Electron Configurations

antimony, and tellurium) are also in the p block. At the left-hand sideand bottom of the block, there are eight p-block metals. The locationsof the nonmetals, metalloids, and metals in the p block are shown withdistinctive colors in Figure 5-6 and in the periodic table printed on theinside back cover of this textbook.

The elements of Group 17 (fluorine, chlorine, bromine, iodine, andastatine) are known as the halogens. The halogens are the most reactivenonmetals. They react vigorously with most metals to form examples ofthe type of compound known as salts. As you will see later, the reactiv-ity of the halogens is based on the presence of seven electrons in theirouter energy levels—one electron short of the stable noble-gas config-uration. Fluorine and chlorine are gases at room temperature, bromineis a reddish liquid, and iodine is a dark purple solid. Astatine is asynthetic element prepared in only very small quantities. Most of itsproperties are estimated, although it is known to be a solid.

The metalloids, or semiconducting elements, fall on both sides of aline separating nonmetals and metals in the p block. They are mostlybrittle solids with some properties of metals and some of nonmetals.The metalloid elements have electrical conductivity intermediatebetween that of metals, which are good conductors, and nonmetals,which are nonconductors.

The metals of the p block are generally harder and denser than thes-block alkaline-earth metals, but softer and less dense than the d-blockmetals.With the exception of bismuth, these metals are sufficiently reac-tive to be found in nature only in the form of compounds. Onceobtained as free metals, however, they are stable in the presence of air.

Fluorine Bromine Iodine

Chlorine

NSTA

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C H A P T E R 5138

The electron configurations of atoms of four elements are written at the top of page 139.For each element, name the block and group in the periodic table in which it is located.Then name the element by consulting the periodic table on pages 130–131. Identify each

SAMPLE PROBLEM 5- 4

The f-Block Elements: Lanthanides and ActinidesIn the periodic table, the f-block elements are wedged between Groups3 and 4 in the sixth and seventh periods. Their position reflects the factthat they involve the filling of the 4f sublevel. With seven 4f orbitals tobe filled with two electrons each, there are a total of 14 f-block elementsbetween lanthanum, La, and hafnium, Hf, in the sixth period. Thelanthanides are shiny metals similar in reactivity to the Group 2alkaline-earth metals.

There are also 14 f-block elements, the actinides, between actinium,Ac, and element 104, Rf, in the seventh period. In these elements the 5fsublevel is being filled with 14 electrons. The actinides are all radioac-tive. The first four actinides (thorium, Th, through neptunium, Np) havebeen found naturally on Earth. The remaining actinides are known onlyas laboratory-made elements.

Without looking at the periodic table, write the outer electron configuration for the Group14 element in the second period. Then name the element, and identify it as a metal, non-metal, or metalloid.

The group number is higher than 12, so the element is in the p block. The total number ofelectrons in the highest occupied s and p sublevels is therefore equal to the group numberminus ten, 14 − 10 = 4. With two electrons in the s sublevel, two electrons must also be pre-sent in the 2p sublevel, giving an outer electron configuration of 2s22p2. The element iscarbon, C, which is a nonmetal.

SOLUTION

SAMPLE PROBLEM 5- 3

Answer1. a. 3s23p5

b. chlorine, nonmetal

2. a. fourth period, p block,Group 15

b. arsenic, metalloid

1. a. Without looking at the periodic table, writethe outer electron configuration for theGroup 17 element in the third period.

b. Name the element described in (a), andidentify it as a metal, nonmetal, or metalloid.

2. a. Without looking at the periodic table,identify the period, block, and group of anelement with the electron configuration[Ar]3d104s24p3.

b. Name the element described in (a), andidentify it as a metal, nonmetal, or metalloid.

PRACTICE



SECTION REVIEW

1. Into what four blocks can the periodic table bedivided to illustrate the relationship between theelements’ electron configurations and their place-ment in the periodic table?

2. What name is given to each of the followinggroups of elements on the periodic table:a. Group 1 c. Groups 3–12 e. Group 18b. Group 2 d. Group 17

3. What are the relationships between group configuration and group number for elements in the s, p, and d blocks?

4. Without looking at the periodic table, write theouter electron configuration for the Group 15 element in the fourth period.

5. Without looking at the periodic table, identify theperiod, block, and group of the element with theelectron configuration [Ar]3d 74s2.

a. The 4f sublevel is filled with 14 electrons. The 5d sublevel is partially filled with nineelectrons. Therefore, this is a d-block element. The element is the transition metal plat-inum, Pt, which is in Group 10 and has a low reactivity.

b. The incompletely filled p sublevel shows this to be a p-block element. With a total ofseven electrons in the ns and np sublevels, this is a Group 17 element, a halogen. The element is chlorine, Cl, and is highly reactive.

c. This element has a noble-gas configuration and thus is in Group 18 in the p block. Theelement is argon, Ar, an unreactive nonmetal and a noble gas.

d. The incomplete 4f sublevel shows that it is an f-block element and a lanthanide. Groupnumbers are not assigned to the f block. The element is samarium, Sm. The lanthanidesare all reactive metals.

SOLUTION

Answer1. a. p block, second period,

Group 17, halogens, fluorine,nonmetal, high reactivity

b. d block, fourth period, Group11, transition elements, cop-per, metal, low reactivity

c. s block, fifth period, Group1, alkali metals, rubidium,metal, high reactivity

1. For each of the following, identify the block,period, group, group name (where appropriate),element name, element type (metal, nonmetal, ormetalloid), and relative reactivity (high or low):

a. [He]2s22p5

b. [Ar]3d104s1

c. [Kr]5s1

PRACTICE

element as a metal, nonmetal, or metalloid. Finally, describe it as likely to be of highreactivity or of low reactivity.

a. [Xe]4f 145d96s1 c. [Ne]3s23p6

b. [Ne]3s23p5 d. [Xe]4f 66s2


Chlorinenucleus

ChlorinenucleusDistance

between nuclei

Atomicradius99 pm

198 pm

ElectronConfiguration andPeriodic Properties

S o far, you have learned that the elements are arranged in the peri-odic table according to their atomic number and that there is a roughcorrelation between the arrangement of the elements and their electronconfigurations. In this section, the relationship between the periodic lawand electron configurations will be further explored.

Atomic Radii

Ideally, the size of an atom is defined by the edge of its orbital. However,this boundary is fuzzy and varies under different conditions. Therefore,to estimate the size of an atom, the conditions under which the atomexists must be specified. One way to express an atom’s radius is to mea-sure the distance between the nuclei of two identical atoms that arechemically bonded together, then divide this distance by two. As illus-trated in Figure 5-12, atomic radius may be defined as one-half the dis-tance between the nuclei of identical atoms that are bonded together.

Period TrendsFigure 5-13 gives the atomic radii of the elements and Figure 5-14 pre-sents this information graphically. Note that there is a gradual decreasein atomic radii across the second period from lithium, Li, to neon, Ne.

C H A P T E R 5140

FIGURE 5-12 One method ofdetermining atomic radius is tomeasure the distance between thenuclei of two identical atoms thatare bonded together in an elementor compound, then divide thisdistance by two. The atomic radiusof a chlorine atom, for example, is99 picometers (pm).

SECTION 5-3

OBJECTIVES

Define atomic and ionic radii,ionization energy, electronaffinity, and electronegativity.

Compare the periodic trendsof atomic radii, ionizationenergy, and electronegativity,and state the reasons forthese variations.

Define valence electrons, andstate how many are presentin atoms of each main-groupelement.

Compare the atomic radii,ionization energies, and electronegativities of the d-block elements with thoseof the main-group elements.


The trend to smaller atoms across a period is caused by the increasingpositive charge of the nucleus. As electrons add to s and p sublevels inthe same main energy level, they are gradually pulled closer to the morehighly charged nucleus.This increased pull results in a decrease in atom-ic radii. The attraction of the nucleus is somewhat offset by repulsionamong the increased number of electrons in the same outer energylevel. As a result, the difference in radii between neighboring atoms ineach period grows smaller, as shown in Figure 5-13.

Group TrendsExamine the atomic radii of the Group 1 elements in Figure 5-13. Noticethat the radii of the elements increase as you read down the group. Aselectrons occupy sublevels in successively higher main energy levelslocated farther from the nucleus, the sizes of the atoms increase. In gen-eral, the atomic radii of the main-group elements increase down a group.

Now examine the radii of the Group 13 elements. Although gallium,Ga, follows aluminum, Al, it has a slightly smaller atomic radius thandoes aluminum. This is because gallium, unlike aluminum, is precededin its period by the 10 d-block elements. The expected increase in gal-lium’s radius caused by the filling of the fourth main-energy level is out-weighed by a shrinking of the electron cloud caused by a nuclear chargethat is considerably higher than that of aluminum.


FIGURE 5-13 Atomic radiidecrease from left to right across a period and increase down a group.

Periodic Table of Atomic Radii (pm)

Module 3: Periodic Properties

CHEMISTRYCHEMISTRY

INTERACTIV

E•

T U T O R

H 1

37

N 7

75

P 15

110

O 8

73

F 9

72

Ne 10

71

He 2

31

S 16

103

Cl 17

100

Ar 18

98

As 33

120

Se 34

119

Br 35

114

Kr 36

112

Sb 51

140

Te 52

142

I 53

133

Xe 54

131

Bi 83

150

Po 84

168

B 5

85

Al 13

143

C 6

77

Si 14

118

Ga 31

135

Ge 32

122

In 49

167

Sn 50

140

Tl 81

170

Pb 82

175

At 85

140

Rn 86

141

Li 3

152

Be 4

112

Na 11

186

Mg 12

160

Sc 21

162

Ti 22

147

V 23

134

Cr 24

128

Y 39

180

Zr 40

160

Nb 41

146

Mo 42

139

K 19

227

Ca 20

197

Rb 37

248

Sr 38

215

Cs 55

265

Ba 56

222

Ta 73

146

W 74

139

Mn 25

127

Tc 43

136

Re 75

137

Fe 26

126

Ru 44

134

Os 76

135

Co 27

125

Rh 45

134

Ir 77

136

Fr 87

270

Ra 88

220

Ni 28

124

Pd 46

137

Pt 78

139

Cu 29

128

Ag 47

144

Au 79

144

Zn 30

134

Cd 48

149

Hg 80

151

Er 68

176

Tm 69

176

Yb 70

—

Lu 71

174

Fm 100

—

Md101

—

Dy 66

178

Ho 67

176

Cf 98

186

Es 99

186

No 102

—

Lr 103

—

Ce 58

182

Pr 59

182

Th 90

179

Pa 91

163

Nd 60

181

U 92

156

Pm 61

183

Np 93

155

Sm 62

180

Pu 94

159

Eu 63

208

Am 95

173

Gd 64

180

Cm 96

174

Tb 65

177

Bk 97

Lanthanide series

Actinide series



Group 18

Group 1 Group 2

La 57

183

Hf 72

159

Ac 89

188

C 6

77

Atomicsymbol

Atomicnumber

Atomic radius Relativeatomic size

Db 105

—

Sg 106

—

Bh 107

—

Hs 108

—

Mt 109

—

Rf 104

—

Peri

od

1

2

3

4

5

6

7

Period

1

2

3

4

5

6

7 115 Uuh 116

—

113 Uuq 114

—

117 Uuo 118

—

Uub 112

—

Uun 110

—

Uuu 111

—

—


Atomic number

Ato

mic

rad

ius

(pm

)

0

50

100

150

200

250

300

0 10 20 30 40 50 60 70 80 90 100

Period 2

Period 3

Period 4

Period 5

Period 6

Period 7

HHe

Li

Ne

Na

Ar

K

Kr

Rb

Xe

Cs

Rn

Fr

Es

Atomic Radius vs. Atomic Number

C H A P T E R 5142

a. Of the elements magnesium, Mg, chlorine, Cl, sodium, Na, and phosphorus, P, which hasthe largest atomic radius? Explain your answer in terms of trends in the periodic table.

b. Of the elements calcium, Ca, beryllium, Be, barium, Ba, and strontium, Sr, which has thelargest atomic radius? Explain your answer in terms of trends in the periodic table.

a. All of the elements are in the third period. Of the four, sodium has the lowest atomicnumber and is the first element in the period. Therefore, sodium has the largest atomicradius because atomic radii decrease across a period.

b. All of the elements are in Group 2. Of the four, barium has the highest atomic numberand is farthest down the group. Therefore, barium has the largest atomic radius becauseatomic radii increase down a group.

SOLUTION

SAMPLE PROBLEM 5- 5

AnswerLi, F

AnswerF, At

1. Of the elements Li, O, C, and F, identify the one with the largest atomicradius and the one with the smallest atomic radius.

2. Of the elements Br, At, F, I, and Cl, identify the one with the smallestatomic radius and the one with the largest atomic radius.

PRACTICE

FIGURE 5-14 The plot of atomic radius versus atomic number shows period and group trends.

Ionization Energy

An electron can be removed from an atom if enough energy is supplied.Using A as a symbol for an atom of any element, the process can beexpressed as follows.

A + energy → A+ + e−

The A+ represents an ion of element A with a single positive charge,referred to as a 1+ ion. An ion is an atom or group of bonded atoms thathas a positive or negative charge. Sodium, for example, forms an Na+ ion.Any process that results in the formation of an ion is referred to asionization.

To compare the ease with which atoms of different elements give upelectrons, chemists compare ionization energies. The energy required toremove one electron from a neutral atom of an element is the ionizationenergy, IE (or first ionization energy, IE1). To avoid the influence of near-by atoms, measurements of ionization energies are made on isolatedatoms in the gas phase. Figure 5-15 gives the first ionization energies forthe elements in kilojoules per mole (kJ/mol). Figure 5-16 presents thisinformation graphically.

T H E P E R I O D I C L A W 143Copyright © by Holt, Rinehart and Winston. All rights reserved.

FIGURE 5-15 In general, firstionization energies increase across a period and decrease down a group.

Periodic Table of Ionization Energies (kJ/mol)

1H

1312

3Li520

11Na496

19K419

37Rb403

55Cs376

87Fr—

4Be900

12Mg

738

20Ca590

38Sr550

56Ba503

88Ra509

5B801

13Al578

31Ga579

49In558

81Tl589

113

6C

1086

14Si787

32Ge762

50Sn709

82Pb716

114Uuq

—

7N

1402

15P

1012

33As947

51Sb834

83Bi703

115

8O

1314

16S

1000

34Se941

52Te869

84Po812

116Uuh

—

9F

1681

17Cl1251

35Br1140

53I

1008

85At—

117

21Sc633

39Y600

57La538

89Ac490

22Ti659

40Zr640

72Hf659

104Rf—

23V651

41Nb652

73Ta761

105Db

—

24Cr653

42Mo

684

74W770

106Sg—

25Mn

717

43Tc702

75Re760

107Bh

—

26Fe762

44Ru710

76Os839

108Hs—

27Co760

45Rh720

77Ir878

109Mt

—

28Ni737

46Pd804

78Pt868

110Uun

—

29Cu746

47Ag731

79Au890

111Uuu

—

30Zn906

48Cd868

80Hg1007

112Uub

—

Group 1

2He2372

10Ne2081

118Ar1521

36Kr1351

54Xe1170

86Rn1038

118Uuo

—

66Dy573

98Cf608

67Ho581

99Es619

68Er589

100Fm627

69Tm597

101Md

635

70Yb603

102No642

58Ce534

90Th587

59Pr527

91Pa570

60Nd533

92U598

61Pm536

93Np600

62Sm545

94Pu585

63Eu547

95Am

578

64Gd592

96Cm

581

65Tb566

97Bk601

71Lu523

103Lr—

Group 18

Group 2



Lanthanide series

Actinide series

6

C1086

Atomicnumber

Symbol

First ionizationenergy

1

2

3

4

5

6

7

Peri

od

1

2

3

4

5

6

7

Period


Atomic number

Firs

t io

niz

atio

n e

ner

gy

(kJ/

mo

l)

0

50

100

150

200

250

300

0 10 20 30 40 50 60 70 80 90 100

First Ionization Energy vs. Atomic Number

Period 5

Period 6

Period 7

Period 2

Period 3

Period 4

H

He

Ne

Ar

Kr

Xe

Rn

No

Ra

CsRbKNaLi

Period TrendsIn Figures 5-15 and 5-16, examine the ionization energies for the firstand last elements in each period. You can see that the Group 1 metalshave the lowest first ionization energies in their respective periods.Therefore, they lose electrons most easily. This ease of electron loss is amajor reason for the high reactivity of the Group 1 (alkali) metals. TheGroup 18 elements, the noble gases, have the highest ionization ener-gies. They do not lose electrons easily. The low reactivity of the noblegases is partly based on this difficulty of electron removal.

In general, ionization energies of the main-group elements increaseacross each period. This increase is caused by increasing nuclear charge.A higher charge more strongly attracts electrons in the same energy level.Increasing nuclear charge is responsible for both increasing ionizationenergy and decreasing radii across the periods. Note that, in general,nonmetals have higher ionization energies than metals do. In each pe-riod, the element of Group 1 has the lowest ionization energy and theelement of Group 18 has the highest ionization energy.

Group TrendsAmong the main-group elements, ionization energies generally decreasedown the groups. Electrons removed from atoms of each succeedingelement in a group are in higher energy levels, farther from the nucleus.Therefore, they are removed more easily. Also, as atomic number

C H A P T E R 5144

FIGURE 5-16 Plot of first ioniza-tion energy, IE1, versus atomic num-ber. As atomic number increases,both the period and the grouptrends become less pronounced.



increases going down a group, more electrons lie between the nucleusand the electrons in the highest occupied energy levels. This partiallyshields the outer electrons from the effect of the nuclear charge.Together, these influences overcome the attraction of the electrons tothe increasing nuclear charge.

Removing Electrons from Positive IonsWith sufficient energy, electrons can be removed from positive ions aswell as from neutral atoms. The energies for removal of additional elec-trons from an atom are referred to as the second ionization energy (IE2),third ionization energy (IE3), and so on.

Table 5-3 shows the first five ionization energies for the elements ofthe first, second, and third periods. You can see that the second ioniza-tion energy is always higher than the first, the third is always higher thanthe second, and so on. This is because as electrons are removed in suc-cessive ionizations, fewer electrons remain within the atom to shield theattractive force of the nucleus. Thus, each successive electron removedfrom an ion feels an increasingly stronger effective nuclear charge (thenuclear charge minus the electron shielding).

The first ionization energies in Table 5-3 show that removing a singleelectron from an atom of a Group 18 element is more difficult thanremoving an electron from atoms of other elements in the same period.This special stability of the noble-gas configuration also applies to ionsthat have noble-gas configurations. Notice in Table 5-3 the largeincreases between the first and second ionization energies of lithium, Li,and between the second and third ionization energies of beryllium, Be.Even larger increases in ionization energy exist between the third and

Period 1 Period 2

H He Li Be B C N O F Ne

IE1 1312 520 900 801 1086 1402 1314 1681 2081

IE2 5250 1757 2427 2353 2856 3388 3374 3952

IE3 11 815 3660 4621 4578 5300 6050 6122

IE4 21 007 6223 7475 7469 8408 9370

IE5 32 827 9445 10 990 11 023 12 178

Period 3

Na Mg Al Si P S Cl Ar

IE1 496 738 578 787 1012 1000 1251 1521

IE2 1451 1817 1577 1903 2251 2297 2666

IE3 6912 2745 3232 2912 3361 3822 3931

IE4 9544 10 540 4356 4957 4564 5158 5771

IE5 13 353 13 628 14 831 6274 7013 6540 7238

4562

7733

11 578

16 091

2372

7298

14 849

25 026

37 830

TABLE 5-3 Ionization Energies (in kJ/mol) for Elements of Periods 1–3


C H A P T E R 5146

Consider two main-group elements A and B. Element A has a first ionization energy of 419 kJ/mol. Element B has a first ionization energy of 1000 kJ/mol. For each element,decide if it is more likely to be in the s block or p block. Which element is more likely toform a positive ion?

Element A has a very low ionization energy, meaning that atoms of A lose electrons easily.Therefore, element A is most likely to be a metal of the s block because ionization energiesincrease across the periods. Element B has a very high ionization energy, meaning thatatoms of B lose electrons with difficulty. We would expect element B to lie at the end of aperiod in the p block. Element A is more likely to form a positive ion because it has a muchlower ionization energy than does element B.

SOLUTION

SAMPLE PROBLEM 5- 6

Answer1. a. Q is in the p block, R is

in the s block, T is in the p block, and X is in the p block.

b. Q and R, and X and T are in the same period. Q and T are in the same group.

c. Q would have the highestionization energy, and Rwould have the lowest.

d. R

e. R

1. Consider the four hypothetical main-groupelements Q, R, T, and X with the outer electronconfigurations indicated below. Then answerthe questions that follow.

Q: 3s23p5 R: 3s1 T: 4d105s25p5 X: 4d105s25p1

a. Identify the block location of each hypothet-ical main-group element.

b. Which of these elements are in the same period? Which are in the same group?

c. Which element would you expect to have thehighest first ionization energy? Which wouldhave the lowest first ionization energy?

d. Which element would you expect to have thehighest second ionization energy?

e. Which of the elements is most likely to forma 1+ ion?

PRACTICE

fourth ionization energies of boron, B, and between the fourth and fifthionization energies of carbon, C. In each case, the jump in ionizationenergy occurs when an ion assumes a noble-gas configuration. Forexample, the removal of one electron from a lithium atom ([He]2s1)leaves the helium noble-gas configuration. The removal of four elec-trons from a carbon atom ([He]2s22p2) also leaves the helium configu-ration. A bigger table would show that this trend continues across theentire periodic system.

Electron Affinity

Neutral atoms can also acquire electrons. The energy change that occurswhen an electron is acquired by a neutral atom is called the atom’s elec-tron affinity. Most atoms release energy when they acquire an electron.

A + e− → A− + energy

In this book, the quantity of energy released is represented by a nega-tive number. On the other hand, some atoms must be “forced” to gainan electron by the addition of energy.

A + e− + energy → A−

The quantity of energy absorbed is represented by a positive number.An ion produced in this way will be unstable and will lose the addedelectron spontaneously.

Figure 5-17 shows the electron affinity in kilojoules per mole for theelements. Figure 5-18 presents these data graphically.

Period TrendsAmong the elements of each period, the halogens (Group 17) gain elec-trons most readily. This is indicated in Figure 5-17 by the large negativevalues of halogens’ electron affinities.The ease with which halogen atomsgain electrons is a major reason for the high reactivities of the Group 17elements. In general, as electrons add to the same p sublevel of atomswith increasing nuclear charge, electron affinities become more nega-tive across each period within the p block. An exception to this trendoccurs between Groups 14 and 15. Compare the electron affinities ofcarbon ([He]2s22p2) and nitrogen ([He]2s22p3). Adding an electron to acarbon atom gives a half-filled p sublevel. This occurs much more easily


FIGURE 5-17 The values listed inparentheses in this periodic table ofelectron affinities are approximate.Electron affinity is estimated to be −50 kJ/mol for each of the lantha-nides and 0 kJ/mol for each of theactinides.

Periodic Table of Electron Affinities (kJ/mol)


CHEMISTRYCHEMISTRY

INTERACTIV

E•

T U T O R

1H

–75.4

3Li

–61.8

11Na–54.8

19K

–50.1

37Rb–48.6

55Cs

– 47.2

87Fr

– 47.0

4Be(0)

12Mg

(0)

20Ca(0)

38Sr(0)

56Ba(0)

88Ra(0)

5B

–27.7

13Al

– 44.1

31Ga–30

49In–30

81Tl–20

113

6C

–126.3

14Si

–138.5

32Ge–135

50Sn–120

82Pb–36

114Uuq

—

7N(0)

15P

–74.6

33As–81

51Sb–107

83Bi

–94.6

115

8O

–146.1

16S

–207.7

34Se

–202.1

52Te

–197.1

84Po–190

116Uuh

—

9F

–339.9

17Cl

–361.7

35Br

–336.5

53I

–305.9

85At–280

117

21Sc

–18.8

39Y

–30.7

57La–50

89Ac—

22Ti–7.9

40Zr

–42.6

72Hf(0)

104Rf—

23V

–52.5

41Nb–89.3

73Ta

–32.2

105Db

—

24Cr

–66.6

42Mo–74.6

74W

–81.5

106Sg—

25Mn

(0)

43Tc–55

75Re–15

107Bh

—

26Fe

–16.3

44Ru–105

76Os–110

108Hs—

27Co–66.1

45Rh

–113.7

77Ir

–156.5

109Mt

—

28Ni

–115.6

46Pd–55.7

78Pt

–212.8

110Uun

—

29Cu

–122.8

47Ag

–130.2

79Au

–230.9

111Uuu

—

30Zn(0)

48Cd

(0)

80Hg

(0)

112Uub

—

Group 1

2He

(0)

10Ne

(0)

18Ar(0)

36Kr(0)

54Xe(0)

86Rn(0)

118Uuo

—

Group 18

Group 2



6

C–126.3

Atomicnumber

Symbol

Electronaffinity

1

2

3

4

5

6

7

Peri

od

1

2

3

4

5

6

7

Period


Atomic number

Elec

tro

n a

ffin

ity

(kJ/

mo

l)0

–100

–200

–300

–4000 10 20 30 40 50 60 70 80 90 100

Electron Affinity vs. Atomic Number

Period 2

Period 3

Period 4

Period 5

Period 6

Period 7

He Ne Ar Kr Xe Rn Ra

HLi Na K Rb Cs Fr

than forcing an electron to pair with another electron in an orbital ofthe already half-filled p sublevel of a nitrogen atom.

Group TrendsTrends for electron affinities within groups are not as regular as trendsfor ionization energies. As a general rule, electrons add with greaterdifficulty down a group. This pattern is a result of two competing fac-tors. The first is a slight increase in effective nuclear charge down agroup, which increases electron affinities. The second is an increase inatomic radius down a group, which decreases electron affinities. In gen-eral, the size effect predominates. But there are exceptions, especiallyamong the heavy transition metals, which tend to be the same size oreven decrease in radius down a group.

Adding Electrons to Negative IonsFor an isolated ion in the gas phase, it is always more difficult to add asecond electron to an already negatively charged ion. Therefore, secondelectron affinities are all positive. Certain p-block nonmetals tend toform negative ions that have noble gas configurations. The halogens doso by adding one electron. For example, chlorine has the configuration[Ne]3s23p5. An atom of chlorine achieves the configuration of the noblegas argon by adding an electron to form the ion Cl− ([Ne]3s23p6).Adding another electron is so difficult that Cl2− never occurs. Atoms of

C H A P T E R 5148

FIGURE 5-18 The plot of electronaffinity versus atomic number showsthat most atoms release energywhen they acquire an electron, asindicated by negative values.



FIGURE 5-19 The ionic radii ofthe ions most common in chemicalcompounds are shown. Cations aresmaller and anions are larger thanthe atoms from which they areformed.

Group 16 elements are present in many compounds as 2− ions. Forexample, oxygen ([He]2s22p4) achieves the configuration of the noblegas neon by adding two electrons to form the ion O2−([He]2s22p6).Nitrogen achieves a neon configuration by adding three electrons toform the ion N3−.

Ionic Radii

Figure 5-19 shows the radii of some of the most common ions of the ele-ments. Positive and negative ions have specific names.

A positive ion is known as a cation. The formation of a cation by theloss of one or more electrons always leads to a decrease in atomic radiusbecause the removal of the highest-energy-level electrons results in asmaller electron cloud. Also, the remaining electrons are drawn closer tothe nucleus by its unbalanced positive charge.

A negative ion is known as an anion. The formation of an anion by theaddition of one or more electrons always leads to an increase in atomicradius. This is because the total positive charge of the nucleus remainsunchanged when an electron is added to an atom or an ion. So the elec-trons are not drawn to the nucleus as strongly as they were before theaddition of the extra electron. The electron cloud also spreads outbecause of greater repulsion between the increased number of electrons.

Period TrendsWithin each period of the periodic table, the metals at the left tend toform cations and the nonmetals at the upper right tend to form anions.Cationic radii decrease across a period because the electron cloudshrinks due to the increasing nuclear charge acting on the electrons in

Periodic Table of Ionic Radii (pm)


CHEMISTRYCHEMISTRY

INTERACTIV

E•

T U T O R

Peri

od

1

2

3

4

5

6

7

1

H– 154

7

N3– 146

15

P3– 212

8

O2– 140

9

F– 133

10

Ne —

2

He —

16

S2– 184

17

Cl– 181

18

Ar —

33

As —

34

Se2– 198

35

Br– 196

36

Kr —

51

Sb3+ 76

52

Te2– 221

53

I– 220

54

Xe —

83

Bi3+ 103

84

Po —

5

B —

13

Al3+ 54

6

C4– 260

14

Si —

31

Ga3+ 62

32

Ge —

49

In3+ 80

50

Sn2+ 118

81

Tl3+ 89

82

Pb2+ 119

85

At —

86

Rn —

3

Li+ 76

4

Be2+ 45

11

Na+ 102

12

Mg2+ 72

21

Sc3+ 75

22

Ti2+ 86

23

V2+ 79

24

Cr2+ 80

39

Y3+ 90

40

Zr —

41

Nb —

42

Mo —

19

K+ 138

20

Ca2+ 100

37

Rb+ 152

38

Sr2+ 118

55

Cs+ 167

56

Ba2+ 136

73

Ta —

74

W —

25

Mn2+ 83

43

Tc —

75

Re —

26

Fe2+ 78

44

Ru —

76

Os —

27

Co2+ 65

45

Rh3+ 67

77

Ir —

87

Fr+ 180

88

Ra2+ 148

28

Ni2+ 69

46

Pd2+ 86

78

Pt2+ 80

29

Cu2+ 73

47

Ag+ 115

79

Au+ 137

30

Zn2+ 74

48

Cd2+ 95

80

Hg2+ 119

57

La3+ 116

72

Hf —

89

Ac3+ 111

6

C4– 260

Atomicnumber

Ionic radius (pm)Ionic symbol

Relativeatomic size

Relativeionic size

105

Db —

106

Sg —

107

Bh —

108

Hs —

109

Mt —

104

Rf —

110

Uun —

111

Uuu —

112

Uub —

113 114

Uuq —

115

—

116

Uuh —

117

—

118

Uuo —

Period

1

2

3

4

5

6

7

Group 1 Group 2



Group 18


C H A P T E R 5150

Group number Group configuration Number of valence electrons

1 ns1 1

2 ns2 2

13 ns2p1 3

14 ns2p2 4

15 ns2p3 5

16 ns2p4 6

17 ns2p5 7

18 ns2p6 8

TABLE 5-4 Valence Electrons in Main-Group Elements

the same main energy level. Starting with Group 15, in which atomsassume stable noble-gas configurations by gaining three electrons,anions are more common than cations. Anionic radii decrease acrosseach period for the elements in Groups 15–18. The reasons for thistrend are the same as the reasons that cationic radii decrease from leftto right across a period.

Group TrendsAs they are in atoms, the outer electrons in both cations and anions arein higher energy levels as one reads down a group. Therefore, just asthere is a gradual increase of atomic radii down a group, there is also agradual increase of ionic radii.

Valence ElectronsChemical compounds form because electrons are lost, gained, orshared between atoms. The electrons that interact in this manner arethose in the highest energy levels. These are the electrons most subjectto the influence of nearby atoms or ions. The electrons available to belost, gained, or shared in the formation of chemical compounds arereferred to as valence electrons. Valence electrons are often located inincompletely filled main-energy levels. For example, the electron lostfrom the 3s sublevel of Na to form Na+ is a valence electron.

For main-group elements, the valence electrons are the electrons inthe outermost s and p sublevels. The inner electrons are in filled energylevels and are held too tightly by the nucleus to be involved in com-pound formation.The Group 1 and Group 2 elements have one and twovalence electrons, respectively, as shown in Table 5-4. The elements ofGroups 13–18 have a number of valence electrons equal to the groupnumber minus 10. In some cases, both the s and p sublevel valence elec-trons of the p-block elements are involved in compound formation. Inother cases, only the electrons from the p sublevel are involved.


Electronegativity

Valence electrons hold atoms together in chemical compounds. In manycompounds, the negative charge of the valence electrons is concentratedcloser to one atom than to another.This uneven concentration of chargehas a significant effect on the chemical properties of a compound. It istherefore useful to have a measure of how strongly one atom attractsthe electrons of another atom within a compound.

Linus Pauling, one of America’s most famous chemists, devised a scaleof numerical values reflecting the tendency of an atom to attract elec-trons. Electronegativity is a measure of the ability of an atom in a chem-ical compound to attract electrons. The most electronegative element,fluorine, is arbitrarily assigned an electronegativity value of four.Valuesfor the other elements are then calculated in relation to this value.

Period TrendsAs shown in Figure 5–20, electronegativities tend to increase across eachperiod, although there are exceptions.The alkali and alkaline-earth metalsare the least electronegative elements. In compounds, their atoms have alow attraction for electrons. Nitrogen, oxygen, and the halogens are themost electronegative elements. Their atoms attract electrons strongly incompounds. Electronegativities tend to either decrease down a group or


FIGURE 5-20 Shown are theelectronegativities of the elementsaccording to the Pauling scale. Themost-electronegative elements arelocated in the upper right of thep block. The least-electronegative elements are located in the lowerleft of the s block.

Periodic Table of Electronegativities

1H2.1

3Li1.0

11Na0.9

19K0.8

37Rb0.8

55Cs0.7

87Fr0.7

4Be1.5

12Mg

1.2

20Ca1.0

38Sr1.0

56Ba0.9

88Ra0.9

5B2.0

13Al1.5

31Ga1.6

49In1.7

81Tl1.8

113

6C2.5

14Si1.8

32Ge1.8

50Sn1.8

82Pb1.8

114Uuq

—

7N3.0

15P2.1

33As2.0

51Sb1.9

83Bi1.9

115

8O3.5

16S2.5

34Se2.4

52Te2.1

84Po2.0

116Uuh

—

9F4.0

17Cl3.0

35Br2.8

53I

2.5

85At2.2

117

21Sc1.3

39Y1.2

57La1.1

89Ac1.1

22Ti1.5

40Zr1.4

72Hf1.3

104Rf—

23V1.6

41Nb

1.6

73Ta1.5

105Db

—

24Cr1.6

42Mo

1.8

74W1.7

106Sg—

25Mn

1.5

43Tc1.9

75Re1.9

107Bh

—

26Fe1.8

44Ru2.2

76Os2.2

108Hs—

27Co1.8

45Rh2.2

77Ir2.2

109Mt

—

28Ni1.8

46Pd2.2

78Pt2.2

110Uun

—

29Cu1.9

47Ag1.9

79Au2.4

111Uuu

—

30Zn1.6

48Cd1.7

80Hg1.9

112Uub

—

Group 1

2He

—

10Ne

—

118Ar—

36Kr3.0

54Xe2.6

86Rn2.4

118Uuo

—

66Dy1.2

98Cf1.3

67Ho1.2

99Es1.3

68Er1.2

100Fm

1.3

69Tm

1.3

101Md

1.3

70Yb1.1

102No

1.3

58Ce1.1

90Th1.3

59Pr1.1

91Pa1.5

60Nd

1.1

92U1.4

61Pm

1.1

93Np

1.4

62Sm

1.2

94Pu1.3

63Eu1.1

95Am

1.3

64Gd

1.2

96Cm

1.3

65Tb1.1

97Bk1.3

71Lu1.3

103Lr—

Group 18

Group 2



Lanthanide series

Actinide series

6

C2.5

Atomicnumber

Symbol

Electronegativity

1

2

3

4

5

6

7

Perio

d

1

2

3

4

5

6

7

Period


Atomic number

Elec

tro

neg

ativ

ity

4.0

3.5

3.0

2.5

2.0

1.5

1.0

0.5

0 10 20 30 40 50 60 70 80 90 100

Period 3

Period 4

Period 5

Period 6

Period 7

Electronegativity vs. Atomic Number

Period 2

H

F

Cl Kr

Xe

Rn

LiNa

K RbCs

remain about the same. The noble gases are unusual in that some of themdo not form compounds and therefore cannot be assigned electronega-tivities. When a noble gas does form a compound, its electronegativity israther high, similar to the values for the halogens.The combination of theperiod and group trends in electronegativity results in the highest valuesbelonging to the elements in the upper right of the periodic table. Thelowest values belong to the elements in the lower left of the table. Thesetrends are shown graphically in Figure 5-21.

C H A P T E R 5152

Among the elements gallium, Ga, bromine, Br, and calcium, Ca, which has the highest elec-tronegativity? Explain why in terms of periodic trends.

The elements are all in the fourth period. Bromine has the highest atomic number and isfarthest to the right in the period. Therefore, it should have the highest electronegativitybecause electronegativity increases across the periods.

1. Consider the five hypothetical main-group elements E, G, J, L, and M with the outerelectron configurations shown at the top of page 153.

SOLUTION

SAMPLE PROBLEM 5- 7

FIGURE 5-21 The plot showselectronegativity versus atomic num-ber for Periods 1–6.

PRACTICE

Periodic Properties of the d- and f-Block Elements

The properties of the d-block elements (which are all metals) vary lessand with less regularity than those of the main-group elements. Thistrend is indicated by the curves in Figures 5-14 and 5-16, which flattenwhere the d-block elements fall in the middle of Periods 4–6.

Recall that atoms of the d-block elements contain from zero to twoelectrons in the s orbital of their highest occupied energy level and oneto ten electrons in the d sublevel of the next-lower energy level.Therefore, electrons in both the ns sublevel and the (n − 1)d sublevelare available to interact with their surroundings.As a result, electrons inthe incompletely filled d sublevels are responsible for many character-istic properties of the d-block elements.

Atomic RadiiThe atomic radii of the d-block elements generally decrease across theperiods. However, this decrease is less than that for the main-groupelements because the electrons added to the (n − 1)d sublevel shield theouter electrons from the nucleus. Also, note in Figure 5-14 that the radiidip to a low and then increase slightly across each of the four periods thatcontain d-block elements. As the number of electrons in the d sublevelincreases, the radii increase because of repulsion among the electrons.

In the sixth period, the f-block elements fall between lanthanum(Group 3) and hafnium (Group 4). Because of the increase in atomicnumber that occurs from lanthanum to hafnium, the atomic radius ofhafnium is actually slightly less than that of zirconium, Zr, the elementimmediately above it. The radii of elements following hafnium in thesixth period vary with increasing atomic number in the usual manner.


Answer1. a. All are in the p block. E, J,

and M are in the same period,and E, G, and L are in thesame group.

b. E should have the highestelectron affinity; E, G, and Lare most likely to form 1−ions; E should have the high-est electronegativity.

c. The ionic radius would belarger.

d. E, G, and L

E = 2s22p5 G = 4d105s25p5 J = 2s22p2

L = 5d106s26p5 M = 2s22p4

a. Identify the block location for each element.Then determine which elements are in the same period and which are in the same group.

b. Which element would you expect to have thehighest electron affinity? Which would youexpect to form a 1− ion? Which should have the highest electronegativity?

c. Compare the ionic radius of the typical ionformed by the element G with the radius of its neutral atom.

d. Which element(s) contains seven valenceelectrons?


Ionization EnergyAs they do for the main-group elements, ionization energies of the d-block and f-block elements generally increase across the periods. Incontrast to the decrease down the main groups, however, the first ion-ization energies of the d-block elements generally increase down eachgroup. This is because the electrons available for ionization in the outers sublevels are less shielded from the increasing nuclear charge by elec-trons in the incomplete (n − 1)d sublevels.

Ion Formation and Ionic RadiiThe order in which electrons are removed from all atoms of the d-blockand f-block elements is exactly the reverse of the order given by theelectron-configuration notation. In other words, electrons in the highestoccupied sublevel are always removed first. For the d-block elements,this means that although newly added electrons occupy the d sublevels,the first electrons to be removed are those in the outermost s sublevels.For example, iron, Fe, has the electron configuration [Ar]3d64s2. It losesa 4s electron first to form Fe+ ([Ar]3d64s1). Fe+ can then lose the second4s electron to form Fe2+ ([Ar]3d6). Fe2+ can then lose a 3d electron toform Fe3+ ([Ar]3d5).

Most d-block elements commonly form 2+ ions in compounds. Some,such as iron and chromium, also commonly form 3+ ions. The Group 3elements form only ions with a 3+ charge. Copper forms 1+ and 2+ ions,and silver usually forms only 1+ ions. As expected, the cations havesmaller radii than the atoms do. Comparing 2+ ions across the periodsshows a decrease in size that parallels the decrease in atomic radii.

ElectronegativityThe d-block elements all have electronegativities between 1.1 and2.54. Only the active metals of Groups 1 and 2 have lower electroneg-ativities. The d-block elements also follow the general trend for electro-negativity values to increase as radii decrease, and vice versa. Thef-block elements all have similar electronegativities, which range from1.1 to 1.5.

C H A P T E R 5154

SECTION REVIEW

1. State the general period and group trends amongmain-group elements with respect to each of thefollowing properties:

a. atomic radii d. ionic radiib. first ionization energy e. electronegativityc. electron affinity

2. Among the main-group elements, what is the rela-tionship between group number and the numberof valence electrons among group members?

3. a. In general, how do the periodic properties ofthe d-block elements compare with those of themain-group elements?

b. Explain the comparisons made in (a).


C H A P T E R 5 R E V I E W

• The periodic law states that the physical andchemical properties of the elements are periodicfunctions of their atomic numbers.

• The periodic table is an arrangement of the ele-ments in order of their atomic numbers so that

elements with similar properties fall in the samecolumn.

• The columns in the periodic table are referred to as groups.

CHAPTER SUMMARY

5-1

actinide (126) lanthanide (126) periodic law (125) periodic table (125)

anion (149) electron affinity (147) ion (143) ionization energy (143)atomic radius (140) electronegativity (151) ionization (143) valence electrons (150)cation (149)

alkali metals (132) halogens (137) main-group elements (136) transition elements (134)alkaline-earth metals (132)

• The rows in the periodic table are called periods.• Many chemical properties of the elements can

be explained by the configurations of the ele-ments’ outermost electrons.

• The noble gases exhibit unique chemical stabilitybecause their highest occupied levels have anoctet of electrons, ns2np6 (with the exception of

helium, whose stability arises from its highestoccupied level being completely filled with twoelectrons, 1s2).

• Based on the electron configurations of the ele-ments, the periodic table can be divided intofour blocks: the s block, the p block, the d block,and the f block.

• The groups and periods of the periodic table dis-play general trends in the following properties ofthe elements: electron affinity, electronegativity,ionization energy, atomic radius, and ionic radius.

• The electrons in an atom that are available to belost, gained, or shared in the formation of chemi-

cal compounds are referred to as valenceelectrons.

• In determining the electron configuration of anion, the order in which electrons are removedfrom the atom is the reverse of the order givenby the atom’s electron-configuration notation.

Vocabulary

Vocabulary

Vocabulary

5-2

5-3

1. Describe the contributions made by the follow-ing scientists to the development of the periodictable:a. Stanislao Cannizzarob. Dmitri Mendeleevc. Henry Moseley (5-1)

2. State the periodic law. (5-1)

3. How is the periodic law demonstrated withinthe groups of the periodic table? (5-1)

REVIEWING CONCEPTS 4. a. How do the electron configurations withinthe same group of elements compare?

b. Why are the noble gases relativelyunreactive? (5-2)

5. What determines the length of each period inthe periodic table? (5-2)

6. What is the relationship between the electronconfiguration of an element and the period inwhich that element appears in the periodictable? (5-2)


CHAPTER 5 REVIEW

C H A P T E R 5156

7. a. What information is provided by the specificblock location of an element?

b. Identify, by number, the groups located within each of the four block areas. (5-2)

8. a. Which elements are designated as the alkalimetals?

b. List four of their characteristic properties. (5-2)

9. a. Which elements are designated as thealkaline-earth metals?

b. How do their characteristic propertiescompare with those of the alkali metals? (5-2)

10. a. Write the usual group configuration notationfor each d-block group.

b. How do the group numbers of those groupsrelate to the number of outer s and delectrons? (5-2)

11. What name is sometimes used to refer to theentire set of d-block elements? (5-2)

12. a. What types of elements make up the p block?b. How do the properties of the p-block

metals compare with those of the metals inthe s and d blocks? (5-2)

13. a. Which elements are designated as thehalogens?

b. List three of their characteristic properties. (5-2)

14. a. Which elements are metalloids?b. Describe their characteristic properties. (5-2)

15. Which elements make up the f block in theperiodic table? (5-2)

16. a. What are the main-group elements?b. What trends can be observed across the

various periods within the main-group elements? (5-2)

17. a. What is meant by atomic radius?b. What trend is observed among the atomic

radii of main-group elements across a period?c. How can this trend be explained? (5-3)

18. a. What trend is observed among the atomicradii of main-group elements down a group?

b. How can this trend be explained? (5-3)

19. Define each of the following terms:a. ion

b. ionizationc. first ionization energyd. second ionization energy (5-3)

20. a. How do the first ionization energies of main-group elements vary across a period anddown a group?

b. Explain the basis for each trend. (5-3)

21. a. What is electron affinity?b. What signs are associated with electron

affinity values, and what is the significance of each sign? (5-3)

22. a. Distinguish between a cation and an anion.b. How does the size of each compare with the

size of the neutral atom from which it isformed? (5-3)

23. a. What are valence electrons?b. Where are such electrons located? (5-3)

24. For each of the following groups, indicatewhether electrons are more likely to be lost orgained in compound formation and give thenumber of such electrons typically involved:a. Group 1 d. Group 16b. Group 2 e. Group 17c. Group 13 f. Group 18 (5-3)

25. a. What is electronegativity?b. Why is fluorine special in terms of electro-

negativity? (5-3)

26. Identify the most- and least-electronegativegroups of elements in the periodic table. (5-3)

Electron Configuration and PeriodicProperties27. Write the noble-gas notation for the electron

configuration of each of the following elements,and indicate the period in which each belongs:a. Li c. Cu e. Snb. O d. Br

28. Without looking at the periodic table, identifythe period, block, and group in which the ele-ments with the following electron configurationsare located. (Hint: See Sample Problem 5-1.)a. [Ne]3s23p4

b. [Kr]4d105s25p2

c. [Xe]4f 145d106s26p5

PROBLEMS


CHAPTER 5 REVIEW


29. Based on the information given below, give thegroup, period, block, and identity of each ele-ment described. (Hint: See Sample Problem 5-2.)a. [He]2s2

b. [Ne]3s1

c. [Kr]5s2

d. [Ar]4s2

e. [Ar]3d54s1

30. Without looking at the periodic table, write theexpected outer electron configuration for eachof the following elements. (Hint: See SampleProblem 5-3.)a. Group 7, fourth periodb. Group 3, fifth periodc. Group 12, sixth period

31. Identify the block, period, group, group name(where appropriate), element name, elementtype, and relative reactivity for the elementswith the following electron configurations.(Hint: See Sample Problem 5-4.)a. [Ne]3s23p1

b. [Ar]3d104s24p6

c. [Kr]4d105s1

d. [Xe]4f 15d16s2

Atomic Radius, Ionization, ElectronAffinity, and Electronegativity32. Of cesium, Cs, hafnium, Hf, and gold, Au, which

element has the smallest atomic radius? Explainyour answer in terms of trends in the periodictable. (Hint: see Sample Problem 5-5.)

33. a. Distinguish between the first, second, andthird ionization energies of an atom.

b. How do the values of successive ionizationenergies compare?

c. Why does this occur?

34. Without looking at the electron affinity table,arrange the following elements in order ofdecreasing electron affinities: C, O, Li, Na, Rb,and F.

35. a. Without looking at the ionization energytable, arrange the following elements in orderof decreasing first ionization energies: Li, O,C, K, Ne, and F.

b. Which of the elements listed in (a) would youexpect to have the highest second ionizationenergy? Why?

36. a. Which of the following cations is least likelyto form: Sr2+, Al3+, K2+?

b. Which of the following anions is least likelyto form: I−, Cl−, O2−?

37. Which element is the most electronegativeamong C, N, O, Br, and S? Which group does itbelong to? (Hint: See Sample Problem 5-7.)

38. The two ions K+ and Ca2+ each have 18 electronssurrounding the nucleus. Which would youexpect to have the smaller radius? Why?

39. Without looking at the periodic table, identifythe period, block, and group in which each ofthe following elements is located:a. [Rn]7s1

b. [Ar]3d24s2

c. [Kr]4d105s1

d. [Xe]4f 145d96s1

40. a. Which elements are designated as the noblegases?

b. What is the most significant property of theseelements?

41. Which of the following does not have a noble-gasconfiguration: Na+, Rb+, O2−, Br−, Ca+, Al3+, S2−?

42. a. How many groups are in the periodic table?b. How many periods are in the periodic table?c. Which two blocks of the periodic table make

up the main-group elements?

43. Write the noble-gas notation for the electronconfiguration of each of the following elements,and indicate the period and group in which eachbelongs:a. Mgb. Pc. Scd. Y

44. Use the periodic table to describe the chemicalproperties of the following elements:a. fluorine, Fb. xenon, Xec. sodium, Nad. gold, Au

MIXED REVIEW


CHAPTER 5 REVIEW

C H A P T E R 5158

45. Identify which trends in the diagrams belowdescribe atomic radius, ionization energy, elec-tron affinity, and electronegativity.

46. For each element listed below, determine thecharge of the ion most likely to be formed andthe identity of the noble gas whose electronconfiguration is thus achieved.a. Li e. Mg i. Brb. Rb f. Al j. Bac. O g. Pd. F h. S

47. Describe some differences between the s-blockmetals and the d-block metals.

48. Why do the halogens readily form 1− ions?

49. The electron configuration of argon differs fromthose of chlorine and potassium by one electron

each. Compare the reactivity of these threeelements.

As a member on the newly inhabited space stationAlpha, you are given the task of organizing informa-tion on newly discovered elements as it comes infrom the laboratory. To date, five elements havebeen discovered and have been assigned names andsymbols from the Greek alphabet. An analysis of thenew elements has yielded the following data:

50. Applying Models Using the data on the fivenew elements, create a periodic table based ontheir properties.

51. Predicting Outcomes Using your newly cre-ated periodic table, predict the atomic numberof an element with an atomic mass of 11.29 thathas nonmetallic properties and is very reactive.

52. Predicting Outcomes Predict the atomic num-ber of an element having an atomic mass of15.02 that exhibits metallic properties but issofter than lambda and harder than beta.

53. Analyzing Information Analyze your periodictable for periodic trends, and describe thetrends that you identify.

CRITICAL THINKING

Element Atomic Atomicname no. mass Properties

Epsilon 23 47.33 nonmetal, very reactive,produces a salt when combined with a metal,gaseous state

Beta 13 27.01 metal, very reactive, softsolid, low melting point

Gamma 12 25.35 nonmetal, gaseous element,extremely unreactive

Delta 4 7.98 nonmetal, very abundant,forms compounds withmost other elements

Lambda 9 16.17 metal, solid state, goodconductor, high luster,hard and dense

Increases

Incr

ease

s

a.

Increases

Incr

ease

s

b.

Increases

Incr

ease

s

c.

ε

β

γ

∆

Λ


CHAPTER 5 REVIEW


57. Prepare a report tracing the evolution of thecurrent periodic table since 1900. Cite thechemists involved and their major contributions.

58. Write a report describing the contributions ofGlenn Seaborg toward the discovery of many ofthe actinide elements.

59. Your teacher will give you an index card thatidentifies the electronegativity, ionization energy,and electron affinity of an element. Identify theelement by analyzing these properties in relationto periodic trends.

60. Construct your own periodic table or obtain aposter that shows related objects, such as fruitsor vegetables, in periodic arrangement. Describethe organization of the table and the trends itillustrates. Use this table to make predictionsabout your subject matter.

ALTERNATIVE ASSESSMENT

RESEARCH & WRITING

HANDBOOK SEARCH

55. Review the boiling point and melting pointdata in the tables of the Elements Handbook(beginning on page 730). Make a list of theelements that exist as liquids or gases at theboiling point of water, 100°C.

56. Because transition metals have vacant dorbitals, they form a greater variety of coloredcompounds than do the metals of Groups 1and 2. Review the section of the ElementsHandbook on transition metals (pages 740–749)and answer the following:

a. What colors are exhibited by chromium inits common oxidation states?

b. What gems contain chromium impurities?c. What colors are often associated with the

following metal ions: copper, cadmium,cobalt, zinc, and nickel?

d. What transition elements are considerednoble metals? What are the characteristicsof a noble metal?

TECHNOLOGY & LEARNING

54. Graphing Calculator Graphing AtomicRadius vs. Atomic Number

The graphing calculator can run a program that graphs data such as atomic radius versusatomic number. Graphing the data within thedifferent periods will allow you to discovertrends. Begin by creating a table of data. Thenuse the program to plot the data.

Go to Appendix C. If you are using a TI 83Plus, you can download the program and datasets and run the application as directed. If youare using another calculator, your teacher willprovide you with keystrokes and data sets touse. Remember that after creating the lists, youwill need to name the program and check thedisplay, as explained in Appendix C. You willthen be ready to run the program. After youhave graphed the data, answer these questions.

a. Would you expect any atomic number tohave an atomic radius of 20? Why or whynot?

b. A relationship is considered a function if itcan pass a vertical line test. That is, if a verti-cal line can be drawn anywhere on the graphand only pass through one point, the rela-tionship is a function. Do either of these setsof data represent a function? If so, whichone(s)?

c. How would you describe the graphical rela-tionship between the atomic numbers andatomic radii?


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