Chapter 2.4 Periodic properties of the elements · properties of the elements. Properties such as atomic radius, ... Aluminum forms an oxide ... further trend from metallic to non-metallic

Chapter 2.4Periodic properties of the elements

Electron Configurations, Atomic Properties, and Periodic Trends

Electron configurations help to determine the atomic and chemical properties of the elements. Properties such as atomic radius, ionization energy, electron affinity and electronegativity are periodic as they follow recurring trends in the periodic table.

As we move down the periodic table the electron configuration becomes longer and more complex. This goes beyond the purpose of our course.

Electron Configurations, Atomic Properties, and Periodic Trends

Periodic Trends in Reactivity of MetalsReaction With oxygen in the air With water Reaction with dilute

acidsGroup 1 React rapidly to form oxides with the

general formula M2O. These compounds are strong bases.

React vigorously with cold water, generating H2(g) and forming a strong base with the general formula MOH.

The reaction is dangerously violent, igniting H2(g) generated.

Group 2 React moderately to form basic oxides with the general formula MO.

React slowly in cold water, generating H2(g) and forming a strong base, M(OH)2.

React quickly generating H2(g).

Group 13 React slowly. Aluminum forms an oxide coating that protects the metal; the compound Al2O3 can act as a base or an acid.

Al displaces H2(g) from steam. Al reacts readily if the protective oxide coating is removed, to generate H2(g) .

Group 14 C and Si are nonmetals, and form acidic oxides. Ge is a metalloid, and forms an acidic oxide. SnO2 can act as an acid or a base. PbO2 is unreactive.

The most metallic elements, Sn and Pb, do not react with water.

Sn and Pb react slowly to form H2(g).

Transition Metals (for example, Fe, Co, Ni, Cu, Zn, Ag, Au)

Zn is the most reactive transition metal, and forms ZnO when the metal is burned in air. Ag and Au do not react.

Zn and Fe displace H2(g)from steam. Fe rusts slowly at room temperature.

Zn reacts readily to generate H2(g), Cu, Ag, and Au do not react.

increase the orbital energy in oxygen. Therefore, less energy is required toremove the fourth p sublevel electron in oxygen.

Periodic trends in ionization energy are linked to trends involving thereactivity of metals. In general, the chemical reactivity of metals increasesdown a group and decreases across a period. These trends, as well as afurther trend from metallic to non-metallic properties across a period, andincreasing metallic properties down a group, are shown in Table 3.1.

Table 3.1 Periodic Trends in Reactivity of Metals

React rapidly to form oxides with the general formula M2O.These compoundsare strong bases.

React vigorouslywith cold water,generating H2(g)and forming astrong base with the general formula MOH.

The reaction isdangerouslyviolent, ignitingH2(g) generated.

ReactionWith oxygenin the air With water

Reaction with dilute acids

Group 1

React moderately to form basic oxides with the general formula MO.

React slowly incold water, generating H2(g)and forming a strong base, M(OH)2.

React quicklygenerating H2(g).

React slowly.Aluminum forms an oxide coating that protects the metal; the compound Al2O3can act as a base or an acid.

Al displaces H2(g)from steam.

Al reacts readilyif the protectiveoxide coating isremoved, to generate H2(g) .

Group 2

Group 13

C and Si are non-metals, and form acidic oxides. Geis a metalloid, and forms an acidic oxide. SnO2 can act as an acid or a base. PbO2 is unreactive.

The most metallicelements, Sn and Pb, do not reactwith water.

Sn and Pb reactslowly to formH2(g) .

Group 14

Zn is the most reactive transitionmetal, and forms ZnOwhen the metal is burned in air. Ag and Au do not react.

Zn and Fe displace H2(g) from steam. Ferusts slowly atroom temperature.

Zn reacts readilyto generate H2(g), Cu, Ag, and Au do not react.

Transition Metals(for example, Fe,Co, Ni, Cu, Zn, Ag, Au)

Chapter 3 Atoms, Electrons, and Periodic Trends • MHR 155

Magnesium ribbon reactingwith oxygen in air

Potassium reacting with water

Nickel reacting with dilute acid

Atomic Radius

The size of an atom, its atomic radius, decreases across a period. Furthermore, atomic radii generally increase down a group. Two factors affect differences in atomic radii:

1.As n increases, there is a higher probability of finding electrons farther from their nucleus. Therefore, the atomic volume is larger.

2.The other factor that affects atomic radii is changing nuclear charge — specifically, the effective nuclear charge (the net force of attraction between electrons and the nucleus)

152 MHR • Unit 2 Structure and Properties

Periodic Trends in Atomic RadiusAs you know, atoms are not solid spheres. Their volumes (the extent oftheir orbitals) are described in terms of probabilities. Nevertheless, thesize of an atom — its atomic radius — is a measurable property. Chemistscan determine it by measuring the distance between the nuclei of bonded,neighbouring atoms. For example, for metals, atomic radius is half the distance between neighbouring nuclei in a crystal of the metal element.For elements that occur as molecules, which is the case for many non-metals, atomic radius is half the distance between nuclei of identicalatoms that bonded together with a single covalent bond. In Figure 3.23,the radii of metallic elements represent the radius of an atom in a metalliccrystal. The radii of all other elements represent the radius of an atom ofthe element participating in a single covalent bond with one additional,like atom.

Representations of atomic radii for main group and transition elements.(Values for atomic radii are given in picometres. Those in parentheses have only two significant digits.)

Figure 3.23

3

4

5

6

7

2

1

1(IA)

3(IIIB)

4(IVB)

5(VB)

6(VIB)

7(VIIB)

8 9 10(VIIIB)

11(IB)

12(IIB)

2(IIA)

3(IIIA)

4(IVA)

5(VA)

6(VIA)

7(VIIA)

8(VIIIA)

4

5

6

H He

Li Be

Na Mg

K Ca

Rb Sr

Cs Ba

Fr Ra

Sc Ti

Y Zr

Hf

V Cr

Nb Mo

Ta W

Mn Fe

Tc Ru

Re Os

Co

Rh

IrLa

Ni Cu

Pd Ag

Pt Au

Zn

Cd

Hg

Ga Ge

In Sn

Tl Pb

As

Sb

Bi

Se Br

Te I

Po At

Kr

Xe

Rn

B C

Al Si

N

P

O F

S Cl

Ne

Ar

37 31

152 112

186 160

227 197

248 215

265 222

(270) (220)

162 147

180 160

159

134 128

146 139

146 139

127 126

136 134

137 135

125

134

136187

124 128

137 144

138 144

134

151

151

135 122

167 140

170 146

120

140

150

119 114

142 133

168 (140)

112

131

(140)

85 77

143 118

75

110

73 72

103 100

71

98

Peri

od

First Ionization Energy

The energy needed to completely remove one electron from a ground state gaseous atom is called the ionization energy. Energy must be added to the atom to remove an electron in order to overcome the force of attraction exerted on the electron by the nucleus. Since multi-electron atoms have two or more electrons, they also have more than one ionization energy.

For calcium the first ionization energy (IE1), is 599 kJ/mol:

The second ionization energy (IE2) is the amount of energy required to remove the second electron. For calcium, it may be represented as:

For a given element, IE2 is always greater than IE1 because it is always more difficult to remove a negatively charged electron from a positively charged ion than from the corresponding neutral atom.

Ionization energies measure how strongly electrons are bound to atoms. Ionization always requires energy to remove an electron from the attractive force of the nucleus. Low ionization energies indicate easy removal of electrons, and hence easy positive ion (cation) formation.

Ca(g) + 599 kJ → Ca+(g) + e-

Ca+(g) + 1145 kJ → Ca2+1(g) + e-

SECTION 7.4 Ionization Energy 261

Ioni

zatio

n en

ergy

(kJ

/mol

)

Increasing ionization energy

Incr

easi

ng io

niza

tion

ener

gy

1A2A

3A4A

5A6A

7A8A

He2372

H1312

Li520

Na496

K419

Rb403

Cs376

Ne2081

Ar1521

Kr1351

Xe1170

Rn1037

Be899

Mg738

Ca590

Sr549Ba

503

Sc633

Y600

Lu524

Ti659

Zr640Hf659

V651

Nb652

Ta761

Cr653

Mo684W

770

Mn717

Tc702

Re760

Fe763

Ru710

Os840

Co760

Rh720

Ni737Pd

804Pt

870

Ir880 Au

890

Ag731

Cu746

Hg1007

Cd868

Zn906

B801

Al578

Ga579

In558

Tl589

C1086

Si786

Ge762

Sn709

Pb716 Bi

703

Sb834

As947

P1012

N1402

Po812

Te869

Se941

S1000

O1314

I1008

Br1140

Cl1251

F1681

! FIGURE 7.9 Trends in first ionization energies of the elements.

G O F I G U R EWhich has a larger first ionization energy, Ar or As? Why?

2p

Oxygen

2p

Nitrogen

! FIGURE 7.10 2p orbital filling innitrogen and oxygen.

G O F I G U R EExplain why it is easier to removea 2p electron from an oxygen atomthan from a nitrogen atom.

The irregularities in a given period are subtle but still readily explained. For exam-ple, the decrease in ionization energy from beryllium ([He]2s2) to boron ([He]2s22p1),shown in Figure 7.9, occurs because the third valence electron of B must occupy the 2psubshell, which is empty for Be. Recall that the 2p subshell is at a higher energy than the2s subshell (Figure 6.24). The decrease in ionization energy when moving from nitrogen([He]2s22p3) to oxygen ([He]2s22p4) is because of the repulsion of paired electrons inthe p4 configuration (" FIGURE 7.10). Remember that according to Hund’s rule, eachelectron in the p3 configuration resides in a different p orbital, which minimizes theelectron–electron repulsion among the three 2p electrons. •(Section 6.8)

SAMPLE EXERCISE 7.6 Periodic Trends in Ionization Energy

Referring to a periodic table, arrange the atoms Ne, Na, P, Ar, K in order of increasing first ion-ization energy.

SOLUTIONAnalyze and Plan We are given the chemical symbols for five elements. To rank them ac-cording to increasing first ionization energy, we need to locate each element in the periodictable. We can then use their relative positions and the trends in first ionization energies to pre-dict their order.

between the electron and the nucleus. As this attraction increases, it becomes more difficultto remove the electron and, thus, the ionization energy increases. As we move across a pe-riod, there is both an increase in effective nuclear charge and a decrease in atomic radius,causing the ionization energy to increase. As we move down a column, the atomic radius in-creases while the effective nuclear charge increases rather gradually. Thus, the attraction be-tween the nucleus and the electron decreases, causing the ionization energy to decrease.

Trends in first ionization energies of the elements.

First Ionization EnergyEvery element exhibits a large increase in ionization energy when an inner electron is removed. This supports the idea that only the outermost electrons are involved in the chemical bondings and reactions. The inner electrons are too tightly bound to the nucleus to be lost from the atom or even shared with another atom.

First ionization energies vs atomic number for the first 38 elements of the P.T.E.The noble gases have very high IE1 while the 1A metals (Li, Na, K and Rb) have low IE1. Note the similarities in the variations for the Period 2 elements to those for the Period 3 elements. Variations for B group elements are not nearly so pronounced as those for A group elements.

180 C H A P T E R 5 • C H E M I C A L P E R I O D I C I T Y

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ionization energies indicate easy removal of electrons, and hence easy positive ion (cation) formation. Figure 5-2 shows a plot of first ionization energy versus atomic number for several elements.

Elements with low ionization energies (IE) easily lose electrons to form cations.

Figure 5-2 shows that each noble gas has the highest first ionization energy in its period. This should not be surprising because the noble gases are known to be very unre-active elements. It requires more energy to remove an electron from a helium atom (slightly less than 4.0 3 10218 J/atom, or 2372 kJ/mol) than to remove one from a neutral atom of any other element.

He(g) 1 2372 kJ h He1(g) 1 e2

The Group 1A metals (Li, Na, K, Rb, Cs) have very low first ionization energies. Each of these elements has only one electron in its outermost shell (. . . ns1), and they are the largest atoms in their periods. The first electron added to a shell is easily removed to form a noble gas configuration. As we move down the group, the first ionization energies become smaller. The force of attraction of the positively charged nucleus for electrons decreases as the square of the distance between them increases. So as atomic radii increase in a given group, first ionization energies decrease because the outermost electrons are farther from the nucleus.

Effective nuclear charge, Zeff, increases going from left to right across a period. The in-crease in effective nuclear charge causes the outermost electrons to be held more tightly, making them harder to remove. The first ionization energies therefore generally increase from left to right across the periodic table. The reason for the trend in first ionization ener-

Figure 5-2 A plot of first ionization energies for the first 38 elements versus atomic number. The noble gases have very high first ionization energies, and the 1A metals have low first ionization energies. Note the similarities in the variations for the Period 2 elements, 3 through 10, to those for the Period 3 elements, 11 through 18, as well as for the later A group elements. Variations for B group elements are not nearly so pronounced as those for A group elements.

2500

2000

1500

1000

500

05 10 15 20 25 30 35

Atomic number

Firs

t ion

izat

ion

ener

gy (k

J/m

ol)

Period 2 Period 3 Period 4

H

He

Li

Be

B

C

N

O

F

Ne

Na

Mg

Al

Si

P

S

Cl

Ar

K

CaSc

Ti

V

CrMn

Fe

Co Cu

Zn

GaGe

AsBr

Kr

Sr

Rb

SeNi

▶ By Coulomb’s Law, F ~(q1 ) (q2 )

d2 ,

the attraction for the outer shell electrons is directly proportional to the effective charges and inversely proportional to the square of the distance between the charges. Even though the effective nuclear charge increases going down a group, the greatly increased size causes a weaker net attraction for the outer electrons and thus results in a lower first ionization energy.

General trends in first ionization energies of A group elements with position in the periodic table. Exceptions occur at Groups 3A and 6A.

First IEDec

reas

e

Increase

H1312

Li520

Na496

K419

Rb403

Cs377

Be899

Mg738

Ca599

Sr550

Ba503

Sc631

Y617

La538

Ti658

Zr661

Hf681

V650

Nb664

Ta761

Cr652

Mo685

W770

Mn717

Tc702

Re760

Fe759

Ru711

Os840

Co758

Rh720

Ir880

Ni757

Pd804

Pt870

Cu745

Ag731

Au890

Zn906

Cd868

Hg1007

B801

Al578

Ga579

In558

Tl589

C1086

Si786

Ge762

Sn709

Pb715

N1402

P1012

As947

Sb834

Bi703

O1314

S1000

Se941

Te869

Po812

F1681

Cl1251

Br1140

I1008

At890

He2372

Ne2081

Ar1521

Kr1351

Xe1170

Rn1037

Table 5-1 First Ionization Energies (kJ/mol of atoms) of Some Elements

10663_05_ch05_p173-206.indd 180 12/7/12 4:56 PM

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First Ionization Energy

The first ionization energy of an atom is a measure of the energy change when removing an electron from the atom to form a cation.

Electron affinity

Ionization energy measures the ease with which an atom loses an electron, whereas electron affinity measures the ease with which an atom gains an electron.

For example, the first ionization energy of Cl(g), 1251 kJ/mol, is the energy change associated with the process

The positive ionization energy means that energy must be put into the atom to remove the electron.

Cl(g) + 1251 kJ → Cl+(g) + e-

[Ne]3s23p5 [Ne]3s23p4

Cl(g) - 349 kJ → Cl-(g) - e-

[Ne]3s23p5 [Ne]3s23p6

Most atoms can also gain electrons to form anions. The energy change that occurs when an electron is added to a gaseous atom is called the electron affinity because it measures the attraction, or affinity, of the atom for the added electron.

For most atoms, energy is released when an electron is added

For example, the addition of an electron to a chlorine atom is accompanied by an energy change of -349 kJ/mol, the negative sign indicating that energy is released during the process

Therefore the electron affinity of Cl is -349 kJ/mol.

For the electron affinity the trends through the P.T.E. are not as evident as they are for ionization energy. The halogens, which miss one electron to complete the p orbital, have the most negative electron affinities. By gaining an electron, a halogen atom forms a stable anion that has a noble-gas configuration.

The addition of an electron to a noble gas requires that the electron reside in a higher-energy orbital that is empty in the atom. Because occupying a higher-energy orbital is energetically unfavorable, the electron affinity is highly positive.

Electron affinity

264 CHAPTER 7 Periodic Properties of the Elements

The fact that the electron affinity is positive means that an electron will not attach itselfto an Ar atom; the ion is unstable and does not form.

! FIGURE 7.11 shows the electron affinities for the s- and p-block elements ofthe first five periods. Notice that the trends are not as evident as they are for ionizationenergy. The halogens, which are one electron shy of a filled p subshell, have the most-

negative electron affinities. By gaining an electron, a halogen atom forms astable anion that has a noble-gas configuration (Equation 7.5). The addi-tion of an electron to a noble gas, however, requires that the electron residein a higher-energy subshell that is empty in the atom (Equation 7.6).Because occupying a higher-energy subshell is energetically unfavorable,the electron affinity is highly positive. The electron affinities of Be and Mgare positive for the same reason; the added electron would reside in a previ-ously empty p subshell that is higher in energy.

The electron affinities of the group 5A elements are also interesting.Because these elements have half-filled p subshells, the added electronmust be put in an orbital that is already occupied, resulting in largerelectron–electron repulsions. Consequently, these elements have electronaffinities that are either positive (N) or less negative than the electronaffinities of their neighbors to the left (P, As, Sb). Recall that in Section 7.4

we saw a discontinuity in the trends for first ionization energy for the same reason.Electron affinities do not change greatly as we move down a group (Figure 7.11).

For F, for instance, the added electron goes into a 2p orbital, for Cl a 3p orbital, for Br a 4porbital, and so forth. As we proceed from F to I, therefore, the average distance betweenthe added electron and the nucleus steadily increases, causing the electron–nucleusattraction to decrease. However, the orbital that holds the outermost electron is increas-ingly spread out, so that as we proceed from F to I, the electron–electron repulsions arealso reduced. As a result, the reduction in the electron–nucleus attraction is counterbal-anced by the reduction in electron–electron repulsions.

G I V E I T S O M E T H O U G H TWhat is the relationship between the value for the first ionization energy of a

ion and the electron affinity of Cl(g)?

7.6 | METALS, NONMETALS, AND METALLOIDSAtomic radii, ionization energies, and electron affinities are properties of individualatoms. With the exception of the noble gases, however, none of the elements exist innature as individual atoms. To get a broader understanding of the properties of ele-ments, we must also examine periodic trends in properties that involve large collectionsof atoms.

The elements can be broadly grouped as metals, nonmetals, and metalloids(" FIGURE 7.12). •(Section 2.5) Some of the distinguishing properties of metalsand nonmetals are summarized in # TABLE 7.3.

In the following sections, we explore some common patterns of reactivity across theperiodic table. We will examine reactivity for nonmetals and metals in more depth inlater chapters.

Cl-(g)

Ar-

H!73

1A

Be" 0

Li!60Na

!53

K!48

Rb!47

Mg" 0

Ca!2

Sr!5

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

8A

B!27

Al!43

Ga!30

In!30

C!122

Si!134

Ge!119

Sn!107

N" 0

P!72

As!78

Sb!103

O!141

S!200

Se!195

Te!190

F!328

Cl!349

Br!325

I!295

He" 0

Ne" 0

Ar" 0

Kr" 0

Xe" 0

$ FIGURE 7.11 Electron affinity inkJ/mol for selected s- and p-blockelements.

G O F I G U R EWhich of the groups shown herehas the most negative electronaffinities? Why does this makesense?

TABLE 7.3 • Characteristic Properties of Metals and Nonmetals

Metals Nonmetals

Have a shiny luster; various colors, although most are silvery Do not have a luster; various colorsSolids are malleable and ductile Solids are usually brittle; some are hard, some are softGood conductors of heat and electricity Poor conductors of heat and electricityMost metal oxides are ionic solids that are basic Most nonmetal oxides are molecular substances that form acidic solutionsTend to form cations in aqueous solution Tend to form anions or oxyanions in aqueous solution

Electron affinity in kJ/mol for some elements

With the exception of the noble gases, however, none of the elements exist in nature as individual atoms. The elements can be broadly grouped, considering the properties, as metals, nonmetals, and metalloids.

Metals, Nonmetals and Metalloids

SECTION 7.6 Metals, Nonmetals, and Metalloids 265

The more an element exhibits the physical and chemical properties of metals, thegreater its metallic character. As indicated in Figure 7.12, metallic character generallyincreases as we proceed down a group of the periodic table and decreases as we proceedright across a period. Let’s now examine the close relationships that exist between elec-tron configurations and the properties of metals, nonmetals, and metalloids.

MetalsMost metallic elements exhibit the shiny luster we associate with metals (! FIGURE7.13). Metals conduct heat and electricity. In general they are malleable (can bepounded into thin sheets) and ductile (can be drawn into wires). All are solids at roomtemperature except mercury , which is a liquid at room tem-perature. Two metals melt at slightly above room temperature, cesium at andgallium at . At the other extreme, many metals melt at very high temperatures.For example, chromium melts at .

Metals tend to have low ionization energies (Figure 7.9) and therefore tend to formcations relatively easily. As a result, metals are oxidized (lose electrons) when they un-dergo chemical reactions. Among the fundamental atomic properties (radius, electronconfiguration, electron affinity, and so forth), first ionization energy is the best indicatorof whether an element behaves as a metal or a nonmetal.

! FIGURE 7.14 shows the oxidation states of representative ions of metals andnonmetals. As noted in Section 2.7, the charge on any alkali metal ion in a compound isalways , and that on any alkaline earth metal is always . For atoms belonging toeither of these groups, the outer s electrons are easily lost, yielding a noble-gas electronconfiguration. For metals belonging to groups with partially occupied p orbitals (groups3A–7A), cations are formed either by losing only the outer p electrons (such as ) orthe outer s and p electrons (such as ). The charge on transition-metal ions doesnot follow an obvious pattern. One characteristic of the transition metals is their abilityto form more than one cation. For example, iron is in some compounds and in others.

G I V E I T S O M E T H O U G H TDescribe a general relationship between trends in metallic character and trendsin ionization energy.

3+2+

Sn4+Sn2+

2+1+

1900 °C29.8 °C

28.4 °C(melting point = -39 °C)

Increasing metallic character

Incr

easi

ng m

etal

lic c

hara

cter

1H

1A1

4Be

3Li

11Na

19K

21Sc

37Rb

55Cs

87Fr

12Mg

20Ca

38Sr

56Ba

88Ra

23V

41Nb

73Ta

105Db

39Y

57La

89Ac

22Ti

40Zr

72Hf

104Rf

2A2

3B3

4B4

5B5

6B6

7B7

3A13

4A14

5A15

6A16

7A17

8A18

1B11

2B128 9

24Cr

42Mo

74W

106Sg

25Mn

43Tc

75Re

5B

13Al

31Ga

49In

81Tl

6C

14Si

32Ge

50Sn

82Pb

7N

15P

33As

51Sb

83Bi

8O

16S

34Se

52Te

84Po

9F

17Cl

35Br

53I

85At

2He

10Ne

18Ar

36Kr

54Xe

86Rn

71Lu

103Lr

70Yb

102No

69Tm

101Md

68Er

100Fm

67Ho

99Es

66Dy

98Cf

65Tb

97Bk

64Gd

96Cm

63Eu

95Am

62Sm

94Pu

61Pm

93Np

60Nd

92U

59Pr

91Pa

58Ce

90Th

26Fe

44Ru

76Os

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

48Cd

80Hg

27Co

10

45Rh

77Ir

8B

107Bh

108Hs

109Mt

110Ds

111Rg

112Cp

114 118116 117113 115

Metals

Metalloids

Nonmetals

" FIGURE 7.12 Metals, metalloids, and nonmetals.

G O F I G U R ENotice that germanium, Ge, is a metalloid but tin, Sn, is a metal. What changesin atomic properties do you think are important in explaining this difference?

" FIGURE 7.13 Metals are shiny andmalleable.

Metals, Nonmetals and Metalloids

Metals Nonmetals•Have a shiny luster; various colors, although most are silvery •Solids are malleable and ductile

•Good conductors of heat and electricity•Most metal oxides are ionic solids that are basic•Tend to form cations in aqueous solution

•Do not have a luster; various colors

•Solids are usually brittle; some are hard, some are soft•Poor conductors of heat and electricity•Most nonmetal oxides are molecular substances that form acidic solutions •Tend to form anions or oxyanions in aqueous solution

Characteristic Properties of Metals and Nonmetals

Most metallic elements exhibit the shiny luster and conduct heat and electricity. In general they are malleable (can be pounded into thin sheets) and ductile (can be drawn into wires). All are solids at room temperature except mercury. Two metals (Cs and Ga ) melt above room temperature. At the other extreme, many metals have very high melting temperatures. For example, chromium melts at 1900 °C.







3+2+

Sn4+Sn2+

2+1+

1900 °C29.8 °C



Incr

easi

ng m

etal

lic c

hara

cter

1H

1A1

4Be

3Li

11Na

19K

21Sc

37Rb

55Cs

87Fr

12Mg

20Ca

38Sr

56Ba

88Ra

23V

41Nb

73Ta

105Db

39Y

57La

89Ac

22Ti

40Zr

72Hf

104Rf

2A2

3B3

4B4

5B5

6B6

7B7

3A13

4A14

5A15

6A16

7A17

8A18

1B11

2B128 9

24Cr

42Mo

74W

106Sg

25Mn

43Tc

75Re

5B

13Al

31Ga

49In

81Tl

6C

14Si

32Ge

50Sn

82Pb

7N

15P

33As

51Sb

83Bi

8O

16S

34Se

52Te

84Po

9F

17Cl

35Br

53I

85At

2He

10Ne

18Ar

36Kr

54Xe

86Rn

71Lu

103Lr

70Yb

102No

69Tm

101Md

68Er

100Fm

67Ho

99Es

66Dy

98Cf

65Tb

97Bk

64Gd

96Cm

63Eu

95Am

62Sm

94Pu

61Pm

93Np

60Nd

92U

59Pr

91Pa

58Ce

90Th

26Fe

44Ru

76Os

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

48Cd

80Hg

27Co

10

45Rh

77Ir

8B

107Bh

108Hs

109Mt

110Ds

111Rg

112Cp

114 118116 117113 115

Metals

Metalloids

Nonmetals










3+2+

Sn4+Sn2+

2+1+

1900 °C29.8 °C



Incr

easi

ng m

etal

lic c

hara

cter

1H

1A1

4Be

3Li

11Na

19K

21Sc

37Rb

55Cs

87Fr

12Mg

20Ca

38Sr

56Ba

88Ra

23V

41Nb

73Ta

105Db

39Y

57La

89Ac

22Ti

40Zr

72Hf

104Rf

2A2

3B3

4B4

5B5

6B6

7B7

3A13

4A14

5A15

6A16

7A17

8A18

1B11

2B128 9

24Cr

42Mo

74W

106Sg

25Mn

43Tc

75Re

5B

13Al

31Ga

49In

81Tl

6C

14Si

32Ge

50Sn

82Pb

7N

15P

33As

51Sb

83Bi

8O

16S

34Se

52Te

84Po

9F

17Cl

35Br

53I

85At

2He

10Ne

18Ar

36Kr

54Xe

86Rn

71Lu

103Lr

70Yb

102No

69Tm

101Md

68Er

100Fm

67Ho

99Es

66Dy

98Cf

65Tb

97Bk

64Gd

96Cm

63Eu

95Am

62Sm

94Pu

61Pm

93Np

60Nd

92U

59Pr

91Pa

58Ce

90Th

26Fe

44Ru

76Os

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

48Cd

80Hg

27Co

10

45Rh

77Ir

8B

107Bh

108Hs

109Mt

110Ds

111Rg

112Cp

114 118116 117113 115

Metals

Metalloids

Nonmetals




Metals are shiny and malleable

Metals tend to have low IE and tend to form cations. The charge on any alkali metal ion in a compound is always 1+, and that on any alkaline earth metal is always 2+, that is, the outer s electrons are easily lost.Compounds made up of a metal and a nonmetal tend to be ionic substances

Among the fundamental atomic properties (radius, electron configuration, etc.), first ionization energy is the best indicator of whether an element behaves as a metal or a nonmetal.

Metals

SECTION 7.7 Trends for Group 1A and Group 2A Metals 269

Group 1A: The Alkali MetalsThe alkali metals are soft metallic solids (! FIGURE 7.19). All have characteristic metallicproperties, such as a silvery, metallic luster and high thermal and electrical conductivity.The name alkali comes from an Arabic word meaning “ashes.”Many compounds of sodiumand potassium, two alkali metals, were isolated from wood ashes by early chemists.

As " TABLE 7.4 shows, the alkali metals have low densities and melting points, andthese properties vary in a fairly regular way with increasing atomic number. We see theusual trends as we move down the group, such as increasing atomic radius and decreas-ing first ionization energy. The alkali metal of any given period has the lowest I1 value inthe period (Figure 7.9), which reflects the relative ease with which its outer s electroncan be removed. As a result, the alkali metals are all very reactive, readily losing one elec-tron to form ions carrying a charge. •(Section 2.7)

The alkali metals exist in nature only as compounds. Sodium and potassium are rela-tively abundant in Earth’s crust, in seawater, and in biological systems, usually as the cationsof ionic compounds. All alkali metals combine directly with most nonmetals. For example,they react with hydrogen to form hydrides and with sulfur to form sulfides:

[7.16]

[7.17]

where M represents any alkali metal. In hydrides of the alkali metals (LiH, NaH, and soforth), hydrogen is present as , the hydride ion. A hydrogen atom that has gained anelectron, this ion is distinct from the hydrogen ion, , formed when a hydrogen atomloses its electron.

The alkali metals react vigorously with water, producing hydrogen gas and a solu-tion of an alkali metal hydroxide:

[7.18]

These reactions are very exothermic. In many cases enough heat is generated to ignite theH2, producing a fire or sometimes even an explosion (# FIGURE 7.20). The reaction ismost violent for the heavier alkali metals, in keeping with their lower ionization energies.

2 M(s) + 2 H2O(l) ¡ 2 MOH(aq) + H2(g)

H+H-

2 M(s) + S(s) ¡ M2S(s)

2 M(s) + H2(g) ¡ 2 MH(s)

1+

" FIGURE 7.19 Sodium, like the otheralkali metals, is soft enough to be cut witha knife.

TABLE 7.4 • Some Properties of the Alkali Metals

ElementElectronConfiguration

Melting Point (°C)

Density(g/cm3)

Atomic Radius (Å)

I1(kJ/mol)

Lithium [He]2s1 181 0.53 1.34 520

Sodium [Ne]3s1 98 0.97 1.54 496

Potassium [Ar]4s1 63 0.86 1.96 419

Rubidium [Kr]5s1 39 1.53 2.11 403

Cesium [Xe]6s1 28 1.88 2.25 376

Li

$ FIGURE 7.20 Thealkali metals reactvigorously with water.

Na K

Elemental Na

Nonmetals can be solid, liquid, or gas. They are not lustrous and generally are poor conductors of heat and electricity. Their melting points are generally low (although diamond, a form of C, melts at 3570 °C). Under ordinary conditions, seven nonmetals exist as diatomic molecules. Five of these are gases (H2, N2, O2, F2, and Cl2), one is a liquid (Br2), and one is a volatile solid (I2). Excluding the noble gases, the remaining nonmetals are solids that can be either hard, such as diamond, or soft, such as sulfur

Nonmetals

Because of their relatively large electron affinities, nonmetals tend to gain electrons and form anions when they react with metals. For example, the reaction of aluminum with bromine produces the ionic compound aluminum bromide:

Sulfur, known to the medieval world as “brimstone,” is a nonmetal.


SAMPLE EXERCISE 7.8 Metal Oxides

(a) Would you expect scandium oxide to be a solid, liquid, or gas at room temperature? (b) Write the balanced chemical equation for the reaction of scandium oxide with nitric acid.

SOLUTIONAnalyze and Plan We are asked about one physical property of scandium oxide—its state atroom temperature—and one chemical property—how it reacts with nitric acid.

Solve(a) Because scandium oxide is the oxide of a metal, we expect it to be an ionic solid. Indeed itis, with the very high melting point of .(b) In compounds, scandium has a charge, , and the oxide ion is . Consequently,the formula of scandium oxide is Sc2O3. Metal oxides tend to be basic and, therefore, to reactwith acids to form a salt plus water. In this case the salt is scandium nitrate, Sc(NO3)3:

PRACTICE EXERCISEWrite the balanced chemical equation for the reaction between copper(II) oxide and sulfuricacid.Answer:

NonmetalsNonmetals can be solid, liquid, or gas. They are not lustrous and generally are poor con-ductors of heat and electricity. Their melting points are generally lower than those ofmetals (although diamond, a form of carbon, is an exception and melts at ).Under ordinary conditions, seven nonmetals exist as diatomic molecules. Five of theseare gases (H2, N2, O2, F2, and Cl2), one is a liquid (Br2), and one is a volatile solid (I2).Excluding the noble gases, the remaining nonmetals are solids that can be either hard,such as diamond, or soft, such as sulfur (! FIGURE 7.16).

Because of their relatively large, negative electron affinities, nonmetals tend to gainelectrons when they react with metals. For example, the reaction of aluminum withbromine produces the ionic compound aluminum bromide:

[7.12]

A nonmetal typically will gain enough electrons to fill itsoutermost occupied p subshell, giving a noble-gas electron con-figuration. For example, the bromine atom gains one electronto fill its 4p subshell:

Compounds composed entirely of nonmetals are typicallymolecular substances that tend to be gases, liquids, or low-melt-ing solids at room temperature. Examples include the commonhydrocarbons we use for fuel (methane, CH4; propane, C3H8;octane, C8H18) and the gases HCl, NH3, and H2S. Many drugsare molecules composed of C, H, N, O, and other nonmetals.For example, the molecular formula for the drug Celebrex isC17H14F3N3O2S. Most nonmetal oxides are acidic, which meansthat those that dissolve in water form acids:

[7.13]

[7.14]

The reaction of carbon dioxide with water (! FIGURE 7.17) accounts for the acidityof carbonated water and, to some extent, rainwater. Because sulfur is present in oil and coal,combustion of these common fuels produces sulfur dioxide and sulfur trioxide. These

P4O10(s) + 6 H2O(l) ¡ 4 H3PO4(aq)

CO2(g) + H2O(l) ¡ H2CO3(aq)

Nonmetal oxide + water ¡ acid

Br ([Ar]4s23d104p5) + e-Q Br- ([Ar]4s23d104p6)

2 Al(s) + 3 Br2(l) ¡ 2 AlBr3(s)

3570 °C

CuO(s) + H2SO4(aq) ¡ CuSO4(aq) + H2O(l)

Sc2O3(s) + 6 HNO3(aq) ¡ 2 Sc(NO3)3(aq) + 3 H2O(l)

O2-Sc3+3+2485 °C

" FIGURE 7.16 Sulfur, known to themedieval world as “brimstone,” is anonmetal.

" FIGURE 7.17 The reaction of CO2with water containing a bromthymol blueindicator. Initially, the blue color tells us thewater is slightly basic. When a piece of solidcarbon dioxide (“dry ice”) is added, the colorchanges to yellow, indicating an acidicsolution. The mist is water dropletscondensed from the air by the cold CO2 gas.







3+2+

Sn4+Sn2+

2+1+

1900 °C29.8 °C



Incr

easi

ng m

etal

lic c

hara

cter

1H

1A1

4Be

3Li

11Na

19K

21Sc

37Rb

55Cs

87Fr

12Mg

20Ca

38Sr

56Ba

88Ra

23V

41Nb

73Ta

105Db

39Y

57La

89Ac

22Ti

40Zr

72Hf

104Rf

2A2

3B3

4B4

5B5

6B6

7B7

3A13

4A14

5A15

6A16

7A17

8A18

1B11

2B128 9

24Cr

42Mo

74W

106Sg

25Mn

43Tc

75Re

5B

13Al

31Ga

49In

81Tl

6C

14Si

32Ge

50Sn

82Pb

7N

15P

33As

51Sb

83Bi

8O

16S

34Se

52Te

84Po

9F

17Cl

35Br

53I

85At

2He

10Ne

18Ar

36Kr

54Xe

86Rn

71Lu

103Lr

70Yb

102No

69Tm

101Md

68Er

100Fm

67Ho

99Es

66Dy

98Cf

65Tb

97Bk

64Gd

96Cm

63Eu

95Am

62Sm

94Pu

61Pm

93Np

60Nd

92U

59Pr

91Pa

58Ce

90Th

26Fe

44Ru

76Os

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

48Cd

80Hg

27Co

10

45Rh

77Ir

8B

107Bh

108Hs

109Mt

110Ds

111Rg

112Cp

114 118116 117113 115

Metals

Metalloids

Nonmetals










3+2+

Sn4+Sn2+

2+1+

1900 °C29.8 °C



Incr

easi

ng m

etal

lic c

hara

cter

1H

1A1

4Be

3Li

11Na

19K

21Sc

37Rb

55Cs

87Fr

12Mg

20Ca

38Sr

56Ba

88Ra

23V

41Nb

73Ta

105Db

39Y

57La

89Ac

22Ti

40Zr

72Hf

104Rf

2A2

3B3

4B4

5B5

6B6

7B7

3A13

4A14

5A15

6A16

7A17

8A18

1B11

2B128 9

24Cr

42Mo

74W

106Sg

25Mn

43Tc

75Re

5B

13Al

31Ga

49In

81Tl

6C

14Si

32Ge

50Sn

82Pb

7N

15P

33As

51Sb

83Bi

8O

16S

34Se

52Te

84Po

9F

17Cl

35Br

53I

85At

2He

10Ne

18Ar

36Kr

54Xe

86Rn

71Lu

103Lr

70Yb

102No

69Tm

101Md

68Er

100Fm

67Ho

99Es

66Dy

98Cf

65Tb

97Bk

64Gd

96Cm

63Eu

95Am

62Sm

94Pu

61Pm

93Np

60Nd

92U

59Pr

91Pa

58Ce

90Th

26Fe

44Ru

76Os

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

48Cd

80Hg

27Co

10

45Rh

77Ir

8B

107Bh

108Hs

109Mt

110Ds

111Rg

112Cp

114 118116 117113 115

Metals

Metalloids

Nonmetals




2 Al + 3 Br2 → 2 AlBr3(s)

Compounds composed entirely of nonmetals are typically molecular substances that tend to be gases, liquids, or low-melting solids at room temperature. E.g. the common hydrocarbons we use for fuel (methane, CH4; propane, C3H8; octane, C8H18) and the gases HCl, NH3, and H2S.

Metalloids have properties intermediate between those of metals and those of nonmetals. They may have some characteristic metallic properties but lack others. E.g. Silicon looks like a metal, but it is brittle rather than malleable and does not conduct heat or electricity nearly as well as metals do

Metalloids

Several metalloids, e.g. Silicon, are electrical semiconductors and are the principal components of circuits and computer chips. One of the reasons is that their electrical conductivity is intermediate between that of metals and that of nonmetals.

268 CHAPTER 7 Periodic Properties of the Elements

substances dissolve in water to produce acid rain, a major pollutant in many parts of theworld. Like acids, most nonmetal oxides dissolve in basic solutions to form a salt plus water:

[7.15]

G I V E I T S O M E T H O U G H TA compound ACl3 (A is an element) has a melting point of . Would youexpect the compound to be molecular or ionic? If you were told that A is eitherscandium or phosphorus, which do you think is the more likely choice?

SAMPLE EXERCISE 7.9 Nonmetal Oxides

Write the balanced chemical equation for the reaction of solid selenium dioxide, SeO2(s), with(a) water, (b) aqueous sodium hydroxide.

SOLUTIONAnalyze and Plan We note that selenium is a nonmetal. We therefore need to write chemicalequations for the reaction of a nonmetal oxide with water and with a base, NaOH. Nonmetaloxides are acidic, reacting with water to form an acid and with bases to form a salt and water.

Solve(a) The reaction between selenium dioxide and water is like that between carbon dioxide andwater (Equation 7.13):

(It does not matter that SeO2 is a solid and CO2 is a gas under ambient conditions; the point isthat both are water-soluble nonmetal oxides.)

(b) The reaction with sodium hydroxide is like the reaction in Equation 7.15:

PRACTICE EXERCISEWrite the balanced chemical equation for the reaction of solid tetraphosphorus hexoxide withwater.

Answer: P4O6(s) + 6 H2O(l) ¡ 4 H3PO3(aq)

SeO2(s) + 2 NaOH(aq) ¡ Na2SeO3(aq) + H2O(l)

SeO2(s) + H2O(l) ¡ H2SeO3(aq)

-112 °C

CO2(g) + 2 NaOH(aq) ¡ Na2CO3(aq) + H2O(l)

Nonmetal oxide + base ¡ salt + water

! FIGURE 7.18 Elemental silicon.Although it looks metallic, silicon, ametalloid, is brittle and a poor thermaland electrical conductor.

MetalloidsMetalloids have properties intermediate between those of metals and those of nonmetals.They may have some characteristic metallic properties but lack others. For example, themetalloid silicon looks like a metal (" FIGURE 7.18), but it is brittle rather than mal-leable and does not conduct heat or electricity nearly as well as metals do. Compounds ofmetalloids can have characteristics of the compounds of metals or nonmetals.

Several metalloids, most notably silicon, are electrical semiconductors and are theprincipal elements used in integrated circuits and computer chips. One of the reasonsmetalloids can be used for integrated circuits is that their electrical conductivity is inter-mediate between that of metals and that of nonmetals. Very pure silicon is an electricalinsulator, but its conductivity can be dramatically increased with the addition of specificimpurities called dopants. This modification provides a mechanism for controlling theelectrical conductivity by controlling the chemical composition. We will return to thispoint in Chapter 12.

7.7 | TRENDS FOR GROUP 1A AND GROUP 2A METALS

As we have seen, elements in a given group possess general similarities. However, trendsalso exist within each group. In this section we use the periodic table and our knowledgeof electron configurations to examine the chemistry of the alkali metals and alkalineearth metals.

Elemental Si looks metallic but is brittle and has low thermal and electrical conductivity.

Very pure silicon is an electrical insulator, but its conductivity can be dramatically increased with the addition of specific impurities called dopants. This modification provides a mechanism for controlling the electrical conductivity by controlling the chemical composition.

Documents

Chapter 2.4 Periodic properties of the elements · properties of the elements. Properties such as atomic radius, ... Aluminum forms an oxide ... further trend from metallic to non-metallic