7.1 Ions 7 - Ms. Lara La Cueva HS Sciencet1lara.weebly.com/uploads/1/6/3/2/1632178/ch7pdf.pdf ·...

Ionic and Metallic Bonding 187

Print• Guided Reading and Study Workbook,

Section 7.1• Core Teaching Resources, Section 7.1

Review• Transparencies, T75–T78• Small-Scale Chemistry Laboratory

Manual, Lab 10

Technology• Interactive Textbook with ChemASAP,

Problem-Solving 7.1, Assessment 7.1• Go Online, Section 7.1

b ca

Connecting to Your World

187

7.1 Ions

Guide for Reading

Key Concepts• How do you find the number of

valence electrons in an atom of a representative element?

• Atoms of which elements tend to gain electrons? Atoms of which elements tend to lose electrons?

• How are cations formed?• How are anions formed?

Vocabularyvalence electrons

electron dot structures

octet rule

halide ions

Reading StrategySummarizing Write a one-paragraph summary of how the octet rule applies to the forma-tion of ions.

Figure 7.1 Group 4A elements include carbon, silicon, and germanium.

This saw blade contains carbon in the form of diamond. Silicon is used in the manufacture of microchips. Germanium is one of the materials used to make thermoscanning goggles.

a

b

c

Valence ElectronsMendeleev used similarities in the properties of elements to organize hisperiodic table. Scientists later learned that all of the elements within eachgroup of the periodic table behave similarly because they have the samenumber of valence electrons. Valence electrons are the electrons in the high-est occupied energy level of an element’s atoms. The number of valence elec-trons largely determines the chemical properties of an element.

The number of valence electrons is related to the group numbers in theperiodic table. To find the number of valence electrons in an atom of arepresentative element, simply look at its group number. For example, theelements of Group 1A (hydrogen, lithium, sodium, potassium, and soforth) all have one valence electron, corresponding to the 1 in 1A. Carbonand silicon, in Group 4A, have four valence electrons. Nitrogen and phos-phorus, in Group 5A, have five valence electrons; and oxygen and sulfur, inGroup 6A, have six. The noble gases (Group 8A) are the only exceptions tothe group-number rule: Helium has two valence electrons, and all of theother noble gases have eight. Figure 7.1 shows some applications of Group4A elements.

Pyrite (FeS2), a common mineral that emits sparks when struck against steel, is often mistaken for gold—hence its nickname, “fool’s gold.” Although certainly not worth its weight

in gold, pyrite can be used as a source of sulfur in the production of sulfuric acid, a common industrial

chemical. Pyrite is an example of a crystalline solid. In crystalline solids, the component particles of the

substance are arranged in an orderly, repeating fashion. In this chapter, you will learn about crys-

talline solids composed of ions that are bonded together. But first you need to understand how ions

form from neutral atoms.

7.1

FOCUSObjectives7.1.1 Determine the number of

valence electrons in an atom of a representative element.

7.1.2 Explain how the octet rule applies to atoms of metallic and nonmetallic elements.

7.1.3 Describe how cations form.7.1.4 Explain how anions form.

Guide for Reading

Build VocabularyGraphic Organizer Have each stu-dent choose a halide ion. Then have students sketch a graphic organizer that shows how a halide ion relates to the other three vocabulary terms.

Reading StrategyPreview Have students skim the head-ings, look at the visuals, and read the boldfaced text to preview the section.

INSTRUCT

Have students study the photograph and read the text that opens the sec-tion. Discuss the ways in which iron and sulfur ions can interact to form the regu-lar pyrite crystal structure.

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Section Resources

188 Chapter 7

Section 7.1 (continued)

Less Proficient ReadersLead a class discussion on writing electron dot structures. Ask, What is the significance of the electrons that are represented by dots? (They are valence electrons.) How are they placed? (symmetrically around the atom to show placement in orbitals according to pairing rules) Why are the non-valence elec-trons not shown? (Generally, they are not available for chemical bonding.)

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Differentiated Instruction

188 Chapter 7

Table 7.1

Electron Dot Structure of Some Group A Elements

Valence electrons are usually the only electrons used in chemicalbonds. Therefore, as a general rule, only the valence electrons are shown inelectron dot structures. Electron dot structures are diagrams that showvalence electrons as dots. Table 7.1 shows electron dot structures for atomsof some Group A elements. Notice that all of the elements within a givengroup (with the exception of helium) have the same number of electrondots in their structures.

Checkpoint What is an electron dot structure?

The Octet RuleYou learned in Chapter 6 that noble gases, such as neon and argon, areunreactive in chemical reactions. That is, they are stable. In 1916, chemistGilbert Lewis used this fact to explain why atoms form certain kinds of ionsand molecules. He called his explanation the octet rule: In forming com-pounds, atoms tend to achieve the electron configuration of a noble gas. Anoctet is a set of eight. Recall that each noble gas (except helium) has eightelectrons in its highest energy level and a general electron configuration ofns 2np6. Thus the octet rule takes its name from this fact about noble gases.

Atoms of the metallic elements tend to lose their valence electrons,leaving a complete octet in the next-lowest energy level. Atoms of somenonmetallic elements tend to gain electrons or to share electrons withanother nonmetallic element to achieve a complete octet. Although there areexceptions, the octet rule applies to atoms in most compounds.

Formation of CationsAn atom is electrically neutral because it has equal numbers of protons andelectrons; an ion forms when an atom or group of atoms loses or gainselectrons. An atom’s loss of valence electrons produces a cation, or apositively charged ion. Note that for metallic elements, the name of a cationis the same as the name of the element. For example, a sodium atom (Na)forms a sodium cation (Na�). Likewise, a calcium atom (Ca) forms a cal-cium cation (Ca2�). Although their names are the same, there are manyimportant chemical differences between metals and their cations. Sodiummetal, for example, reacts explosively with water. By contrast, sodium cat-ions are quite unreactive. As you may know, they are a component of tablesalt, a compound that is very stable in water.

Word OriginsOctet comes from the Greek word okto, meaning “eight.” There are eight electrons in the highest occupied energy level of the noble gases, except for helium. How do you think the term octetmight also be applied to music or poetry?

Group

Period 1A 2A 3A 4A 5A 6A 7A 8A

1

2

3

4

H He

Li Be B C N O F Ne

Na Mg Al Si P S Cl Ar

K Ca Ga Ge As Se Br Kr

Valence ElectronsUse VisualsTable 7.1 Have students identify the total number of electrons and the number of valence electrons in selected elements from Table 7.1. Reemphasize that the group number equals the number of valence electrons in an atom of a representative element.

TEACHER DemoTEACHER Demo

Valence ElectronsPurpose To model the valence elec-trons of an atom.

Materials plastic egg, 11 marbles

Procedure Hold up a plastic egg con-taining 10 marbles. State that the egg represents a sodium atom and that the marbles represent the 10 electrons making up the n = 1 and n = 2 energy levels. Explain that these electrons can-not be “removed” without “breaking” the egg. Now hold up one additional marble next to the egg. State that this marble represents the eleventh elec-tron, and occupies the n = 3 energy level. This is the valence electron. Explain that if this electron is lost, the resulting atom has an overall 1+ charge.

Expected Outcome Students should be able to distinguish valence electrons from nonvalence electrons.

The Octet RuleDiscussHave students determine the accuracy of this statement: All stable ions of ele-ments result in electronic configura-tions that are isoelectronic with noble gases. (Most of the time this statement is true, but there are exceptions. Use Cu(I) as an example. Explain that a noble gas configuration is not generally possible with elements that would have to gain or lose many electrons to become stable.)

Word OriginsIn music, an octet refers to a group of eight performers, or a composi-tion written for eight musicians. In poetry, an octet refers to a group of eight lines of verse.

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Formation of CationsUse VisualsFigure 7.2 Display the equation that goes with Figure 7.2 on an overhead projector. Point out that the interaction between atoms that produces bond-ing involves only the outermost elec-trons of the atoms. The inner electrons are locked tightly in filled energy levels and do not participate in bonding. Use a colored pen to circle the outermost electron in the sodium atom in this equation. Remind students that the outermost electrons are called valence electrons. Use a different colored pen to circle the octet of electrons in the sodium ion’s highest energy level. Cir-cle the corresponding octet of elec-trons in neon to show the similarity in electron configurations. Have students draw a similar diagram for calcium.

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Answers to...Figure 7.2 8 electrons

Checkpoint

A diagram that shows valence electrons as dots.

Gifted and TalentedHave students show how the atoms of tran-sition elements become stable with pseudo-noble-gas configurations. These atoms would have to gain or lose too many elec-trons to achieve a noble-gas electron config-uration. For example, show how Ag, Zn, and Ga lose 1, 2, and 3 electrons, respectively to form pseudo-noble-gas configurations.

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The most common cations are those produced by the loss of valenceelectrons from metal atoms. Most of these atoms have one to three valenceelectrons, which are easily removed. Sodium, in Group 1A of the periodictable, is typical. Sodium atoms have a total of eleven electrons, includingone valence electron. A sodium atom can lose an electron to become a pos-itively charged sodium ion. The sodium ion has an electron configurationthat is identical to the noble gas neon. When forming a compound, asodium atom loses its one valence electron and is left with an octet (eightelectrons) in what is now its highest energy level. Because the number ofprotons in the sodium nucleus is still eleven, the loss of one unit of negativecharge (an electron) produces a cation with a charge of 1�. You can repre-sent the electron loss, or ionization, of the sodium atom by drawing thecomplete electron configuration of the atom and of the ion formed.

Notice that the electron configuration of the sodium ion (1s22s22p6) is thesame as that of a neon atom. The diagrams below help illustrate this point.

Both the sodium ion and the neon atom have eight electrons in their val-ance shells. Using electron dot structures, you can show the ionizationmore simply.

Na 1s22s22p63s1 ¡ Na 1s22s2 2p6

(1)1*octet

-e-+

Sodium atom

Na

Sodium ion

Na�

Ener

gy le

vel

Neon atom

Ne

2s

1s

2p

Loss of valence electron

3s

Sodium atom Sodium ion electron(electrically (plus sign (minus sign

neutral, indicates one indicatescharge � 0) unit of positive one unit of

charge) negative charge)

Na. ¡ Na+ + e-loss of valence

electron

ionization

Figure 7.2 The sodium atoms in a sodium-vapor lamp ionize to form sodium cations (Na�).Applying Concepts How many electrons are in the highest energy level of Na+?

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For: Links on SodiumVisit: www.SciLinks.orgWeb Code: cdn-1071

190 Chapter 7


DiscussOne way to determine the number of valence electrons in an atom is to look at the electron configuration of the atom. Explain that any electron in an atom outside the noble-gas core is called a valence electron. Using dia-grams such as those on pages 189 and 190, show students several examples of how various atoms of the represen-tative elements form ions and gain a noble-gas electron configuration. Indi-cate the noble-gas core and valence electrons in your diagrams. Lay pieces of magnesium, zinc, copper, and alumi-num on a dry surface in the lab to show that metals do not spontane-ously form metal cations.

CLASS ActivityCLASS

Forming CationsPurpose Students model the formation of cations using equations.

Materials paper, pencil

Procedure Have students write equa-tions similar to that for Mg on student page 190, showing the formation of metal cations from metal atoms. Students should show the electron dot structures for the metal atoms and metal cations that are formed. In addition, you may want students to write out the electron configurations for the metal atoms and cations. Check students’ equations to be sure the correct metal ion is formed.

Expected Outcome Students should be able to use electron dot structures to correctly write equations describing the ionization of metal atoms.

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Special NeedsAsk students to work with a partner to prac-tice drawing electron dot structures for Group A elements. Have one student ran-domly choose an element and write the sym-bol for that element on a piece of paper. Have another student fill in the electron dots for this element. Students can check their work by referring to Table 7.1. Make sure stu-

dents understand that the electron dots rep-resent only valence electrons, not the total number of electrons, and that valence elec-trons are the only electrons involved in chemical reactivity. Also, reinforce the stable octet dot structures for the noble gases and explain how ions of other elements will react to obtain this configuration.

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190 Chapter 7

By losing its lone 4s electron, copper attains a pseudo noble-gas electronconfiguration. Likewise, cations of gold (Au�), cadmium (Cd2�), and mer-cury (Hg2�) also have pseudo noble-gas configurations.

Copper atomCu

Ener

gy le

vel

Copper(l) ionCu�

2s

1s

2p

3s

3p

4s

3d

Loss of valence electron

Li+

Be2+

Na+

Mg2+

K+

Ca2+

Rb+

Sr2+

Cs+

Ba2+

Fr+

Ra2+

1A 2A

Figure 7.3 Walnuts are a good dietary source of magnesium. Magnesium ions (Mg2�) aid in digestive processes.

Figure 7.4 Cations of Group 1A elements have a charge of 1�.Cations of Group 2A elements have a charge of 2�.

Magnesium (atomic number 12) belongs to Group 2A of the periodictable, so it has two valence electrons. A magnesium atom attains the elec-tron configuration of neon by losing both valence electrons. The loss of thevalence electrons produces a magnesium cation with a charge of 2�.

Figure 7.4 lists the symbols of cations formed by metals in Groups 1Aand 2A. Cations of Group 1A elements always have a charge of 1�. Simi-larly, the cations of Group 2A elements always have a charge of 2�. Thisconsistency can be explained in terms of the loss of valence electrons bymetal atoms: The atoms lose enough electrons to attain the electron con-figuration of a noble gas. For example, all Group 2A elements have twovalence electrons. In losing these two electrons, they form 2� cations.

For transition metals, the charges of cations may vary. An atom of iron,for example, may lose two or three electrons. In the first case, it forms theFe2� ion. In the second case, it forms the Fe3� ion.

Some ions formed by transition metals do not have noble-gas electronconfigurations (ns2np6) and are therefore exceptions to the octet rule. Silver,with the electron configuration of 1s22s22p63s23p63d104s24p64d105s1, is anexample. To achieve the structure of krypton, which is the preceding noblegas, a silver atom would have to lose eleven electrons. To acquire the elec-tron configuration of xenon, which is the following noble gas, silver wouldhave to gain seven electrons. Ions with charges of three or greater areuncommon, and these possibilities are extremely unlikely. Thus silver doesnot achieve a noble-gas configuration. But if it loses its 5s1 electron, the con-figuration that results (4s24p64d10 ), with 18 electrons in the outer energy leveland all of the orbitals filled, is relatively favorable in compounds. Such aconfiguration is known as a pseudo noble-gas electron configuration. Silverforms a positive ion (Ag�) in this way. Other elements that behave similarlyto silver are found at the right of the transition metal block. For example, acopper atom can ionize to form a 1� cation (Cu�), as illustrated below.

Magnesium atom Magnesium ion (2 in front(electrically (2� indicates of e� indicates

neutral, two units two units ofcharge � 0) of positive negative charge)

charge)

.Mg. ¡ Mg2+ + 2e-


Formation of AnionsUse VisualsFigure 7.5 Explain that the elements arsenic (As) and tellurium (Te) are met-alloids, not nonmetals. However, they form anions that are named according to the same convention as nonmetal anions are named (arsenide, telluride).

DiscussExplain to students that when one atom forms an ion, that ion is called a monatomic ion. Explain that the prefix mon- or mono- means “one.” Ask stu-dents to predict what ions might be called if they contained more than one atom. Tell students that they will study such ions in Chapter 8.

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Answers to...Figure 7.5 Group 7A

Checkpoint

argon

Water PurificationChlorine gas is often used to purify drinking water; it kills a variety of microorganisms, including those that carry diseases. But chlorine also reacts with organic substances in the water to produce chlorinated compounds such as chloroacetonitrile. Chloroacetonitrile has been shown to cause inflammation of the digestive tract in laboratory animals. As an

alternative to chlorine, some countries have begun purifying water with ozone. Ozone kills microorganisms even more effectively than chlorine. About one percent of the water supply in the United States is now purified with ozone. It is estimated that it would cost $6 billion to switch completely to ozone for treating all the drinking water supplies.

Facts and Figures

Section 7.1 Ions 191

Formation of AnionsAn anion is an atom or a group of atoms with a negative charge. Thegain of negatively charged electrons by a neutral atom produces an anion.Note that the name of an anion of a nonmetallic element is not the same asthe element name. The name of the anion typically ends in -ide. Thus achlorine atom (Cl) forms a chloride ion (Cl�), and an oxygen atom (O)forms an oxide ion (O2�). Figure 7.5 shows the symbols of anions formed bysome elements in Groups 5A, 6A, and 7A.

Because they have relatively full valance shells, atoms of nonmetallicelements attain noble-gas electron configurations more easily by gainingelectrons than by losing them. For example, chlorine belongs to Group 7A(the halogen family) and has seven valence electrons. A gain of one electrongives chlorine an octet and converts a chlorine atom into a chloride ion.

The chloride ion is an anion with a single negative charge. Notice that it hasthe same electron configuration as the noble gas argon.

Chlorine atoms, therefore, need one more valence electron to achieve theelectron configuration of the nearest noble gas. The diagrams below illus-trate how both the chloride ion and the argon atom have an octet of elec-trons in their highest energy levels.

Based on the diagrams above, you use electron dot structures to writean equation showing the formation of a chloride ion from a chlorine atom.

In this equation, each dot in the electron dot structure represents an elec-tron in the valence shell in the electron configuration diagram.

Checkpoint Which noble gas has the same electron configuration as a chloride ion?

Cl 1s22s22p63s23p5 ¡ Cl- 1s22s2 2p63s23p6(1)1*

octet

+e-

Ar 1s22s2 2p63s23p6(1)1*

octet

Chlorine atom

Cl

Chloride ion

Cl

Argon atom

Ar�

Ener

gy le

vel

2s

1s

2p

3s

3p

Gain of valence electron

e

N3–

O2–

P3–

S2–

As3–

Se2–

Te2–

5A 6A

F–

Cl–

Br–

–

7A

Figure 7.5 Atoms of nonmetallic elements form anions by gaining enough valence electrons so as to attain the electron config-uration of the nearest noble gas. Interpreting Diagrams Inwhich group of the periodic table do the elements bromine and iodine belong?

192 Chapter 7


Gifted and TalentedMany common anions are polyatomic ions. Have students infer what a polyatomic ion is, then confirm their definitions. Have them list some common polyatomic anions, such as sulfate (SO4

2−), nitrate (NO3−), and phos-

phate (PO43−). Tell them that polyatomic

ions form ionic compounds and that poly-atomic ions will be studied in Chapter 8.

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DiscussDiscuss with students how certain ele-ments can form either anions or cations. Have students write the electron config-uration of nitrogen (1s22s22p3). Ask, How can a nitrogen atom form a cation that has the electron configu-ration of a noble gas? (It can lose five electrons.) How can a nitrogen atom form an anion that has the electron configuration of a noble gas? (It can gain three electrons.) Have them repeat this process for carbon. Explain that atoms that have few electrons to gain or lose to achieve an octet are more likely to form ions. Other atoms are not likely to form ions, and the process they undergo will be studied in Chapter 8.

CONCEPTUAL PROBLEM 7.1

Answers1. a. sulfide ion, S2–

b. aluminum ion, Al3+

2. a. 2 electrons lost b. 3 electrons gainedc. 2 electrons lost

Practice Problems PlusGive the name and symbol of the ion formed when:a. a nitrogen atom gains three

electrons (nitride ion, N3–)b. a calcium atom loses two

electrons (calcium ion, Ca2+)c. a fluorine atom gains one electron

(fluoride ion, F–)

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192 Chapter 7

The ions that are produced when atoms of chlorine and other halogensgain electrons are called halide ions. All halogen atoms have seven valenceelectrons and need to gain only one electron to achieve the electron config-uration of a noble gas. Thus all halide ions (F�, Cl�, Br�, and I�) have acharge of 1�. The seawater in Figure 7.6 contains many different ions, butthe negatively charged ions—the anions—are mostly chloride ions.

Look at another example. Oxygen is in Group 6A, and oxygen atomseach have six valence electrons. Oxygen atoms attain the electron configu-ration of neon by gaining two electrons, as shown in the diagrams below.

The resulting oxide ions have charges of 2� and are written as O2�. Usingelectron dot structures, you can write the equation for the formation ofoxide ions as follows.

Ener

gy le

vel

Oxygen atom

O

Oxide ion

O

Neon atom

Ne2�

Gain of two valence electrons

2s

1s

2p

Table 7.2 lists some common anions that you will be learning about inthis book. Note that not all of the anions listed end with the suffix -ide.

Checkpoint How many electrons do halogen atoms need to gain in order to achieve the electron configuration of a noble gas?

O � 2e i O 2� �

Figure 7.6 The six most abundant ions in seawater are chloride (Cl�), sulfate (SO4

2�),sodium (Na�), magnesium (Mg2�), calcium (Ca2�), and potassium (K�).

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Table 7.2

Some Common Anions

1� 2� 3�

F� fluoride O2� oxide N3� nitride

Cl� chloride S2� sulfide P3� phosphide

Br� bromide SO42� sulfate PO4

3� phosphate

I� iodide CO32� carbonate

OH� hydroxide

ClO� hypochlorite

NO3� nitrate

C2H3O2� acetate

HCO3� hydrogen

carbonate


Section 7.1 Assessment3. look up the group number of that element4. Atoms of nonmetallic elements tend to

gain electrons; atoms of metallic ele-ments tend to lose electrons.

5. when an atom loses valence electrons6. when an atom gains valence electrons7. a. 1 b. 4 c. 2 d. 6

8. a. b. c. d.

9. a. lose 2 b. gain 1 c. lose 3 d. gain 210. a. potassium cation, K+

b. zinc cation, Zn2+

c. fluoride anion, F−

11. Cd2+: 1s22s22p63s23p63d104s24p64d10

K C Mg O

ASSESSEvaluate UnderstandingHave students refer to the periodic table on page162. To determine the students’ knowledge about the forma-tion of elemental anions and cations, ask the students to determine whether the following ions are likely to exist and why: H− (yes, isoelectronic with He); H+(yes, but without electrons, there is no comparable noble-gas configuration); Sr2+ (yes, isoelectronic with Kr); Al3+ (yes, isoelectronic with Ne); Xe− (no, cannot form ions easily due to stable electron configuration); Zn6− (no, isoelectronic with Kr but formation would require a gain of too many electrons); Zn2+ (yes, not isoelectronic with a noble gas but has pseudo-noble-gas electron configu-ration with eighteen electrons filling up the outer energy level: 3s 23p63d10)

ReteachSelect groups from the periodic table in random order and ask students to predict the common ions that could be formed from elements of each group. Note that predicting is fairly easy for groups at the far left or far right of the table, but more difficult for groups in the center of the table, which have partially filled d and f orbitals.

Connecting Concepts

The amount of energy needed to remove the one valence electron of a Group 1A metal atom (first ioniza-tion energy) is low. After this elec-tron is lost, the outer energy level contains an octet and is stable. If it loses another electron, it will be less stable, so the second ionization energy is high.

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Checkpoint

1 electron

Section 7.1 Ions 193

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Writing the Symbols and Names of IonsThe beaker shown on the right contains iodine vapor. Write the symbol and name of the ion formed whena. an iodine atom gains one electron.b. a strontium atom loses two electrons.

Analyze Identify the relevant concepts.

a. An atom that gains electrons forms a nega-tively charged ion (anion). The name of an anion of a nonmetallic element ends in -ide.

b. An atom that loses electrons forms a positive-ly charged ion (cation). The name of a cation of a metallic element is the same as the name of the element.

Solve Apply concepts to this situation.

a. I�, iodide ion (an anion)b. Sr2�, strontium ion (a cation)

1. Write the name and symbol of the ion formed whena. a sulfur atom gains two electrons.b. an aluminum atom loses three electrons.

2. How many electrons are lost or gained in forming each ion?a. Ba2� b. As3� c. Cu2�

7.1 Section Assessment

3. Key Concept How can you determine the number of valence electrons in an atom of a rep-resentative element?

4. Key Concept Atoms of which elements tend to gain electrons? Atoms of which elements tend to lose electrons?

5. Key Concept How do cations form?

6. Key Concept How do anions form?

7. How many valence electrons are in each atom? a. potassium b. carbon c. magnesium d. oxygen

8. Write the electron dot structure for each element in Question 7.

9. How many electrons will each element gain or lose in forming an ion?

a. calcium (Ca) b. fluorine (F) c. aluminum (Al) d. oxygen (O)

10. Write the name and symbol of the ion formed when a. a potassium atom loses one electron. b. a zinc atom loses two electrons. c. a fluorine atom gains one electron.

11. Write the electron configuration of Cd2�.

Problem-Solving 7.1 Solve Problem 1 with the help of an interactive guided tutorial.

Ionization Energy Reread Section 6.3. How does the octet rule explain the large increase in energy between the first and second ionization energies of Group 1A metals?

Assessment 7.1 Test yourself on the concepts in Section 7.1.

Practice Problems

194 Chapter 7



Review, Interpreting Graphics• Transparencies, T79–T81• Probeware Laboratory Manual,

Section 7.2


Animation 8, Simulation 5, Problem-Solving 7.12, Assessment 7.2

• Go Online, Section 7.2

7.2

FOCUSObjectives7.2.1 Explain the electrical charge of

an ionic compound.7.2.2 Describe three properties of

ionic compounds.

Guide for Reading

Build VocabularyParaphase Have students skim through the section to locate the mean-ings of the vocabulary terms. Then have them paraphrase each definition.

Reading StrategyIdentify Main Ideas/Details Have students identify the main idea of each paragraph of this section. Have them include these ideas in an outline of the section.

INSTRUCT

Have students examine the section-opening photograph and Figure 7.8. Dis-cuss how the reactive (and poisonous) elements sodium metal and chlorine gas can combine to form harmless table salt. Ask, What characteristics of sodium and chlorine atoms allow them to form the stable compound sodium chloride, also known as table salt? (Sodium atoms can lose an electron easily, and chlorine atoms can accept an electron easily. The resulting ions can combine with the other oppositely charged ions.) Remind students that NaCl is an example of an ionic compound.

Formation of Ionic CompoundsDiscussExplain that the formation of positive ions and of negative ions are simulta-neous and interdependent processes. An ionic compound is the result of the transfer of electrons from one set of atoms to another set of atoms. An ionic compound consists entirely of ions.

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Section Resources

194 Chapter 7

You have heard of harvesting crops such as wheat or rice—but salt? In many coastal countries that have warm, relatively dry climates, salt is produced by the evaporation of seawater. The salty water is chan-neled into a series of shallow ponds, where it becomes more concentrated as the water evaporates by exposure to the sun. When the saltwater is concentrated enough, it is diverted into a pan, on which the sodium chloride crystals deposit. Salt farmers then drain the pans and collect the salt into piles to dry. In this section, you will learn how cations and anions combine to form stable compounds such as sodium chloride.

7.2 Ionic Bonds and Ionic Compounds

Guide for Reading

Key Concepts• What is the electrical charge of

an ionic compound?• What are three properties of

ionic compounds?

Vocabularyionic compounds

ionic bonds

chemical formula

formula unit

coordination number

Reading StrategyPreviewing Before you read this section, rewrite the headings as how, why, and what questionsabout ionic compounds. As you read, write answers to the questions.

Formation of Ionic CompoundsCompounds composed of cations and anions are called ionic compounds.Ionic compounds are usually composed of metal cations and nonmetalanions. For example, sodium chloride, or table salt, is composed of sodiumcations and chloride anions. Although they are composed of ions, ioniccompounds are electrically neutral. The total positive charge of the cationsequals the total negative charge of the anions.

Ionic Bonds Anions and cations have opposite charges and attract oneanother by means of electrostatic forces. The electrostatic forces that holdions together in ionic compounds are called ionic bonds.

Sodium chloride provides a simple example of how ionic bonds areformed. Consider the reaction between a sodium atom and a chlorineatom. Sodium has a single valence electron that it can easily lose. (If thesodium atom loses its valence electron, it achieves the stable electron con-figuration of neon.) Chlorine has seven valence electrons and can easilygain one. (If the chlorine atom gains a valence electron, it achieves the sta-ble electron configuration of argon.) When sodium and chlorine react toform a compound, the sodium atom gives its one valence electron to achlorine atom. Thus sodium and chlorine atoms combine in a one-to-oneratio and both ions have stable octets.

Na

1

Cl

s 2s 2p s s ps3 s p2 2 6 1 1 2 2 3 32 2 6 2 5 s ps s p1 2 2 32 2 6 2 6

s ps s p1 2 2 3 3

3

2 2 6 2 61s 2s 2p2 2 6

1s 2s 2p2 2 6

Na

Ne Ar

octet

....

. ... Cl....

....

(1)1*octet

(1)1*

octet(1)1*

octet(1)1*

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Animation 8 Take anatomic-level look at the formation of KCl.


DiscussAsk, Why are crystalline ionic com-pounds generally so rigid and brit-tle? (The crystal is rigid because it is held together by a specific three-dimensional array of relatively strong attractive forces between anions and cations, which is accompanied by minimal charge repul-sion of like ions. The crystal is brittle because the attractive interactions are specifically arranged within the crystal structure. If this arrangement is dis-turbed, as it would be if the crystal were hit with a hammer, charge repulsion between ions of the same charge can force the crystal to fragment.)

DiscussTo assess students’ prior knowledge about ionic bonds and crystals, ask, What is an ionic bond? (an electical attraction between ions of opposite charge) How do a polyatomic ion and a monatomic ion differ? (A mona-tomic ion is an ion formed from a single atom; a polyatomic ion is a stable unit of two or more tightly bound atoms that carries a charge.) Why are crystals of different ionic compounds different shapes? (The shapes reflect different geometric arrangements of anions and cations with different sizes and charges.)

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Answers to...Figure 7.8 Each ion in the arrange-ment is strongly attracted to its neighbors and repulsions are minimized.

Checkpoint

MgCl2

English LearnersEncourage students to look up and define terms used to describe ionic compounds. Students should define terms such as crystal and formula unit in English and their native language. Emphasize how their strong bonding arrangement accounts for crystals having unique properties.

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Section 7.2 Ionic Bonds and Ionic Compounds 195

Structure of sodium ion and chloride ion

Crystals of sodium chlorideArrangement of Na� ionsand Cl� ions in a crystal

of sodium chloride

Chloride ion (Cl�)

18e�

17p�

18n0

11p�

12n0

10e�

Sodium ion (Na�)

Figure 7.7 shows aluminum and bromine reacting to form the com-pound aluminum bromide. Each aluminum atom has three valence elec-trons to lose. Each bromine atom has seven valence electrons and readilygains one additional electron. Therefore, when aluminum and brominereact, three bromine atoms combine with each aluminum atom.

Formula Units The ionic compound sodium chloride is composed ofequal numbers of sodium cations (Na�) and chloride anions (Cl�). As youcan see in Figure 7.8, the ions in solid sodium chloride are arranged inan orderly pattern. There are no single discrete units, only a continuousarray of ions.

Chemists represent the composition of substances by writing chemicalformulas. A chemical formula shows the kinds and numbers of atoms in thesmallest representative unit of a substance. NaCl, for example, is the chem-ical formula for sodium chloride. Note, however, that the formula NaCldoes not represent a single discrete unit. Because an ionic compound existsas a collection of positively and negatively charged ions arranged in repeat-ing patterns, its chemical formula refers to a ratio known as a formula unit.A formula unit is the lowest whole-number ratio of ions in an ionic com-pound. For sodium chloride, the lowest whole-number ratio of the ions is1:1 (one Na� to each Cl�). Thus the formula unit for sodium chloride isNaCl. Although ionic charges are used to derive the correct formula, theyare not shown when you write the formula unit of the compound.

The ionic compound magnesium chloride contains magnesium cat-ions (Mg2�) and chloride anions (Cl�). In magnesium chloride, the ratio ofmagnesium cations to chloride anions is 1:2 (one Mg2� to two Cl�). So itsformula unit is MgCl2. Because there are twice as many chloride anions(each with a 1� charge) as magnesium cations (each with a 2� charge), thecompound is electrically neutral. In aluminum bromide, described earlier,the ratio of aluminum cations to bromide ions is 1:3 (one Al3� to three Br�

ions), so the formula unit is AlBr3.

Checkpoint What is the formula unit for magnesium chloride?

Figure 7.7 Aluminum metal and the nonmetal bromine react to form an ionic solid, aluminum bromide.

Aluminum bromide (AlBr3)

Bromine (Br2) Aluminum (Al)

Figure 7.8 Sodium cations and chloride anions form a repeating three-dimensional array in sodium chloride (NaCl).Inferring How does the arrangement of ions in a sodium chloride crystal help explain why the compound is so stable?

196 Chapter 7



Answers12. a. KI

b. Al2O313. CaCl2

Practice Problems PlusUse electron dot structures to deter-mine chemical formulas of the ionic compounds formed when the follow-ing elements combine: a. magnesium and chlorine (MgCl2)b. aluminum and sulfur (Al2S3)

Properties of Ionic Compounds

CLASS ActivityCLASS

“Hardness” of WaterPurpose Students detect the presence of ions in water samples.

Materials tap water samples, 10-mL graduated cylinder, potassium thiocy-anate solution, dilute ethanoic acid, sodium oxalate solution, dropper

Procedure Water “hardness” is based on ions present in the water. Have stu-dents bring water samples from home to test for hardness. Test 2-mL samples as follows. Add three drops of potas-sium thiocyanate (KSCN) to the first sample. Add three drops of dilute etha-noic acid, CH3COOH, and three drops of sodium oxalate, Na2C2O4, to the sec-ond sample. Mix well.

Expected Outcome A red color from the iron(III) thiocyanate ion, Fe(SCN)2+, indicates the presence of Fe3+ ions. A white precipitate of calcium oxalate, CaC2O4, indicates the presence of Ca2+ ions.

Download a worksheet on Ionic Compounds to complete, and find additional teacher support on NSTA SciLinks.

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Do You Have “Soft” Water?Explain that water hardness varies with loca-tion and source. Generally, water from groundwater sources is harder than water from surface sources. In the United States, most northeastern, southern, and north

western states have predominantly soft water. Generally, hard water of varying degrees is found in the southwestern and midwestern states.

Facts and Figures

196 Chapter 7

Practice Problems

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Predicting Formulas of Ionic CompoundsThe ionic compound formed from potassium and oxygen is used in ceramic glazes. Use electron dot structures to predict the formulas of the ionic compounds formed from the following elements.a. potassium and oxygen

Analyze Identify the relevant concepts.

Atoms of metals lose their valence electrons when forming an ionic compound. Atoms of nonmetals gain electrons. Enough atoms of each element must be used in the formula so that electrons lost equals electrons gained.

Solve Apply concepts to this situation.

a. Start with the atoms.

In order to have a completely filled valence shell, oxygen must gain two electrons. These electrons come from two potassium atoms, each of which loses one electron.

K O.....

.. and

KO.

2

K

.

.. .. ..�

K.

K. ....

..O

Electrons lost now equals electrons gained. The formula of the compound formed (potas-sium oxide) is K2O.

b. Start with the atoms.

Each nitrogen needs three electrons to have an octet, but each magnesium can lose only two electrons. Thus three magnesium atoms are needed for every two nitrogen atoms.

The formula of the compound formed (mag-nesium nitride) is Mg3N2.

Mg. and.

N..

. ..

2

�

Mg.N

.

Mg..

Mg.. N.

.. ..

N..

. ..Mg

2Mg

2Mg

3.... ....

N 3.... ....

12. Use electron dot structures to determine formu-las of the ionic compounds formed whena. potassium reacts with iodine.b. aluminum reacts with oxygen.

13. What is the formula of the ionic compound com-posed of calcium cations and chloride anions?

Problem-Solving 7.12 Solve Problem 12 with the help of an interactive guided tutorial.

Properties of Ionic CompoundsFigure 7.9 shows the striking beauty of the crystals of some ioniccompounds. Most ionic compounds are crystalline solids at roomtemperature. The component ions in such crystals are arranged in repeat-ing three-dimensional patterns. The composition of a crystal of sodiumchloride is typical. In solid NaCl, each sodium ion is surrounded by sixchloride ions, and each chloride ion is surrounded by six sodium ions. Inthis arrangement, each ion is attracted strongly to each of its neighbors andrepulsions are minimized. The large attractive forces result in a very stablestructure. This is reflected in the fact that NaCl has a melting point ofabout 800°C. Ionic compounds generally have high melting points.

b. magnesium and nitrogen

For: Links on Ionic Compounds

Visit: www.SciLinks.orgWeb Code: cdn-1072


CLASS ActivityCLASS

Types of Ionic CompoundsPurpose Students investigate proper-ties of different classes of ionic com-pounds.

Materials library or Internet access

Procedure Divide the class into groups. Have each group choose a dif-ferent class of ionic compounds to research and write about. For example, one group could work with oxides while another group worked with sul-fides. Initially, each student should work alone to discover information such as where the compounds occur in nature, how they are produced, their physical and chemical properties, and any important uses. Finally, students in each group can pool their information to prepare a class display or report.

Expected Outcome Students will discover that different classes of ionic compounds share some properties with other ionic compounds and have some unique properties.

CLASS ActivityCLASS

Crystal StructuresPurpose Students observe different types of crystals.

Materials crystals of ionic compounds, watch glasses, magnifying glasses

Safety Use only nontoxic crystals. Remind students to not touch the crystals.

Procedure Pass around crystals of ionic compounds of various types in watch glasses. Have the students examine the crystals with magnifying glasses and write down their observa-tions. Make a list of these observations and then discuss them in terms of the underlying ionic lattice structures.

Expected Outcome Students should observe the different geometries of different crystal structures.

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Gifted and TalentedHave students research how minerals are categorized according to their ionic nature. Suggest that their written report include infor-mation concerning the physical properties of

minerals and how these properties are used in mineral identification. Encourage students to include drawings, photos, or examples ofminerals from each category.

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Fluorite (CaF2) Grossularite (Ca3 Al2(SiO4)3)

Aragonite (CaCO3) Barite (BaSO4) and calcite (CaCO3) Wulfenite (PbMoO4)

Beryl (BeAl2(SiO3)6) Franklinite ((Zn, Mn2�, Fe2�)(Fe3�, Mn3�)) Pyrite (FeS2)

Hematite (Fe2O3) Rutile (TiO2) Cinnabar (HgS)

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Simulation 5 Simulate the formation of ionic compounds at the atomic level.

Figure 7.9 The beauty of crystalline solids, such as these, comes from the orderly arrangement of their component ions.

198 Chapter 7


Quick LABQuick LAB

Solutions Containing IonsObjective After completing this activ-ity, students will be able to:

• show that ions in solution conduct an electric current.

Students may think that the solutions contain only one ion. Clarify that each solution contains both cations and anions.

Skills Focus Observing, experimenting, concluding

Prep Time 20 minutesMaterials D-cell batteries, masking tape, 30-cm lengths of bell wire with ends scraped bare, clear plastic cups, distilled water, tap water, vinegar, sucrose, sodium chloride, baking soda, conductivity probe (optional)

Class Time 30 minutes

Safety Students should handle wires with caution. The wires may become hot during the activity.

Expected Outcome When ions are present in solution, the solution con-ducts an electric current.

Analyze and Conclude1. Solutions of vinegar, sodium chloride, and baking soda (and maybe tap water) contain ions and therefore conduct elec-tric current and produce bubbles.

2. Distilled water and sugar solution (and maybe tap water) do not contain ions and therefore do not conduct an electric current or produce bubbles.

3. Answers will vary but should indi-cate that a larger number of batteries will increase the current, which will, in turn, cause the rate at which the bub-bles appear to increase.

For EnrichmentAsk, What gases form the bubbles you observe? (hydrogen and oxygen) What is the source of these gases? (water) Have students collect the gases

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produced and check them for the pres-ence of hydrogen and oxygen. Hydro-gen will burn when a lit splint is placed in it, and a glowing split will begin to flame when placed in oxygen.

198 Chapter 7

The coordination number of an ion is the number of ions of oppositecharge that surround the ion in a crystal. Figure 7.10a shows the three-dimensional arrangement of ions in NaCl. Because each Na� ion is sur-rounded by six Cl� ions, Na� has a coordination number of 6. Each Cl� ionis surrounded by six Na� ions and also has a coordination number of 6.Cesium chloride (CsCl) has a formula unit that is similar to that of NaCl. AsFigure 7.10b illustrates, both compounds have cubic crystals, but theirinternal crystal structures are different. Each Cs� ion is surrounded byeight Cl� ions, and each Cl� ion is surrounded by eight Cs� ions. The anionand cation in cesium chloride each have a coordination number of 8.

Figure 7.10c shows the crystalline form of titanium dioxide (TiO2), alsoknown as rutile. In this compound, the coordination number for the cation(Ti4�) is 6. Each Ti4� ion is surrounded by six O2� ions. The coordinationnumber of the anion (O2�) is 3. Each O2� ion is surrounded by three Ti4� ions.

Another characteristic property of ionic compounds has to do withconductivity. Ionic compounds can conduct an electric current whenmelted or dissolved in water. As Figure 7.11 shows, when sodium chloride ismelted, the orderly crystal structure breaks down. If a voltage is appliedacross this molten mass, cations migrate freely to one electrode and anionsmigrate to the other. This ion movement allows electricity to flow betweenthe electrodes through an external wire. For a similar reason, ionic com-pounds also conduct electricity if they are dissolved in water. When dis-solved, the ions are free to move about in the aqueous solution.

Checkpoint What is the coordination number of Ti4 + in TiO2?

Figure 7.10 Sodium chloride and cesium chloride form cubic crystals. In NaCl, each ion has a coordination number of 6.

In CsCl, each ion has a coordination number of 8.

Titanium dioxide forms tetragonal crystals. In TiO2, each Ti4� ion has a coordination number of 6, while each O2� ion has a coordination number of 3.

a

b

c

Figure 7.11 When sodium chloride melts, the sodium and chloride ions are free to move throughout the molten salt. If a voltage is applied, positive sodium ions move to the negative electrode (the cathode), and negative chloride ions move to the positive electrode (the anode). Predicting What would happen if the voltage was applied across a solution of NaCl dissolved in water?

Cl�

Na�

Cl�

Cs�

Ti4�

O2�

Cesium chloride (CsCl)

Sodium chloride (NaCl)

Rutile (TiO2)

a

c

b

Flow of electrons

Power sourceCurrent meter

Molten salt(801°C—1412°C)Inert metal

electrode(cathode) Inert metal

electrode(anode)

Flow ofelectrons

Cl�

Na�

��

Gifted and TalentedHave students write the formulas for ionic compounds formed from selected pairs of cations and anions. Include various poly-atomic cations and anions as well. Examples:1) K+ and S2−(K2S)2) Ca2+ and HCO3−(Ca(HCO3)2)

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DiscussPoint out that the rusting of iron is the production of iron oxide from iron metal and oxygen gas. Point out that Fe3+ is a stable cation of Fe. (Students should know that O2− is the stable anion of O.) The compound formed from these ions is iron(III) oxide, or Fe2O3.

ASSESSEvaluate UnderstandingName ionic compounds and ask stu-dents to identify the cation, anion, and ratio of cations to anions in each compound.

ReteachEmphasize that an ionic solid is a col-lection of independent ions. There is no joining of individual particles to form molecules. Each ion “belongs” as much to one of its nearest neighbors as it belongs to any other. The arrange-ment in an ionic crystal is such that each ion is surrounded by ions of opposite charge, which produces a strong bonding force.

Elements Handbook

Na+ 1s22s22p6

K+ 1s22s22p63s23p6

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If your class subscribes to the Interactive Textbook, use it to review key concepts in Section 7.2.

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Answers to...Figure 7.11 The solution would conduct an electric current; the Na+ ions in solution would migrate to the negative electrode, and the Cl– ions in solution would migrate to the positive electrode.

Checkpoint

Ti4+ has a coor-dination number of 6 in TiO2.

Section 7.2 Assessment14. electrically neutral15. usually solids at room temperature; have

high melting points; conduct an electric current when melted or dissolved in water

16. electrostatic forces that hold ions together in an ionic compound

17. by writing its chemical formula18. a. K2S b. CaO c. Na2O d. AlN19. a. BaCl2 b. MgO c. Li2O d. CaF2

20. b and d21. Acceptable answers should describe a

solid containing positive sodium ions and negative chloride ions in an alternating, regular, and repeating three-dimensional pattern.

22. The ions are free to move.


Quick LABQuick LAB

Solutions Containing Ions

PurposeTo show that ions in solu-tion conduct an electric current.

Materials

• 3 D-cell batteries

• masking tape

• 2 30-cm lengths of bell wire with ends scraped bare

• clear plastic cup

• distilled water

• tap water

• vinegar

• sucrose

• sodium chloride

• baking soda

• conductivity probe (optional)

Procedure

Probe version available in the Probeware Lab Manual.

1. Tape the batteries together so the pos-itive end of one touches the negative end of another. Tape the bare end of one wire to the positive terminal of the battery assembly and the bare end of the other wire to the negative terminal.CAUTION Bare wire ends can be sharp and scratch skin. Handle with care.

2. Half fill the cup with distilled water. Hold the bare ends of the wires close together in the water. Look for the pro-duction of bubbles. They are a sign that the solution conducts electricity.

3. Repeat Step 2 with tap water, vinegar, and concentrated solutions of sucrose, sodium chloride, and baking soda (sodium hydrogen carbonate).

Analyze and Conclude

1. Which solutions produced bubbles of gas? Explain.

2. Which samples did not produce bub-bles of gas? Explain.

3. Would you expect the same results if you used only one battery? If you used six batteries? Explain your answer.


14. Key Concept How can you describe the elec-trical charge of an ionic compound?

15. Key Concept What properties characterize ionic compounds?

16. Define an ionic bond.

17. How can you represent the composition of an ionic compound?

18. Write the correct chemical formula for the com-pounds formed from each pair of ions.

a. K�, S2� b. Ca2�, O2�

c. Na�, O2� d. Al3�, N3�

19. Write formulas for each compound. a. barium chloride b. magnesium oxide c. lithium oxide d. calcium fluoride

20. Which pairs of elements are likely to form ionic compounds?

a. Cl, Br b. Li, Cl c. K, He d. I, Na

21. Describe the arrangement of sodium ions and chloride ions in a crystal of sodium chloride.

22. Why do ionic compounds conduct electricity when they are melted or dissolved in water?

Handbook

Restoring Electrolytes Read the feature on elec-trolytes on page R8. Write electron configurations for the two principal ions found in body fluids.

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200 Chapter 7

Small-ScaleLAB

Small-ScaleLAB

Analysis of Anions and Cations

ObjectiveAfter completing this activity, students will be able to:• develop tests for various ions.• use the tests to analyze unknown

substances.

Skills Focus Observing, drawing conclusions, designing experiments

Prep Time 40 minutes

Materials pencil, paper, ruler, reaction surface, medicine droppers, pipet, staples or solid Fe, AgNO3, HCl, Pb(NO3)2, Na2SO4, HNO3, Na3PO4, NaOH, KSCN, KI, CaCl2, FeCl3.

Advance Prep See below:

Class Time 40 minutes

Expected Outcome In Figure A, Pb2+ ions form a white precipitate with SO4

2− and PO43−. Ag+ ions form a col-

ored precipitate with SO42− and PO4

3−. The NO3

− ion and the HCl and Fe pro-duce an orange-brown color. In Figure B, KSCN and the Fe3+ ion form a prod-uct that is red. Ca2+ and OH− form a white precipitate.

Analyze1. An intermediate compound, FeCl2,

forms, which reacts with the nitrate ion. An orange-brown color forms.

2. Each of the following pairs of ions produces a visible product that can be used to identify the ion in ques-tion: SO4

2− and Ag+, NO3− and HCl +

Fe, PO43− and Pb2+, Ca2+ and OH−,

Fe3+ and SCN3−, K+ and no solution.3. No, neither of the solutions pro-

duced a visible product.

Solution Preparation0.05M AgNO3 2.1 g in 250 mL0.2M Pb(NO3)2 16.6 g in 250 mL0.2M Na2SO4 7.1 g in 250 mL0.1M Na3PO4 9.5 g Na3PO4•12H2O

in 250 mL0.5M NaOH 20.0 g in 1.0 L0.1M KSCN 2.4 g in 250 mL0.1M KI 4.2 g in 250 mL0.5M CaCl2 13.9 g in 250 mL0.1M FeCl3 6.8 g FeCl3•6H2O

in 25 mL of 1.0M NaCl; dilute to 250 mL

1.0M HCl 82 mL of 12M in 1.0 L1.0M HNO3 63 mL of 15.8M in 1.0 L

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You’re the Chemist1. Mix a drop of each unknown with one drop

of each solution from figure A. Compare the results to those with known solutions.

2. Mix a drop of each unknown with one drop of both solutions from Figure B. Compare the results to those with known solutions.

3. Mix an unknown with one drop of each solu-tion from Figures A and B and compare results.

For EnrichmentHave students hypothesize about why none of the tests in the lab identified the potassium ion. Ask, Why do you think potassium ions don’t form precipitates with any of the anions used in the lab? (All potassium compounds are soluble in water.) Have students read through the Quick Lab on page 142. Have them use this information to design an experiment that will provide a positive test for the potassium ion. The flame test for potassium produces a violet color.

L3

200 Chapter 7

Small-ScaleLAB

Small-ScaleLAB

Analysis of Anions and Cations

PurposeTo develop tests for various ions and use the tests to analyze unknown substances.

Materials

• pencil

• ruler

• medicine droppers

• chemicals shown in Figures A and B

• paper

• reaction surface

• pipet

Procedure On one sheet of paper, draw grids similar to Figure A and Figure B. Draw similar grids on a second sheet of paper. Make each square 2 cm on each side. Place a reaction surface over the grids on one of the sheets of paper and add one drop of each solution or one piece of each solid as shown in Figures A and B. Stir each solution by blowing air through an empty pipet. Use the grids on the second sheet of paper as a data table to record your observations for each solution.

AnalyzeUsing your experimental data, record the answers to the following questions below your data table.

1. Carefully examine the reaction of Fe(s) and HCl in the presence of HNO3. What is unique about this reaction? How can you use it to identify nitrate ion?

2. Which solutions from Figure A are the best for identify-ing each anion? Which solutions from Figure B are the best for identifying each cation? Explain.

3. Can your experiments identify K� ions? Explain.

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

1. Design It! Obtain a set of unknown anion solutions from your teacher and design and carry out a series of tests that will identify each anion.

2. Design It! Obtain a set of unknown cation solutions from your teacher and design and carry out a series of tests that will identify each cation.

3. Design It! Obtain a set of unknown solid ionic com-pounds from your teacher. Design and carry out a series of tests that will identify each ion present.

Na2 4

( 4�)

HNO3

(NO3�)

Na3 4

( 4�)

AgNO3

HCl plus

1 piece

of Fe(s)

Pb(NO3)2

Figure A

Anion Analysis

NaOH

KSCN

KI

(K�)

CaCl2

(Ca2�)

FeCl3

(Fe3�)

Figure B

Cation Analysis




Review• Transparencies, T82–T84• Laboratory Manual, Lab 10


Animation 9, Assessment 7.3

7.3

FOCUSObjectives7.3.1 Model the valence electrons of

metal atoms.7.3.2 Describe the arrangement of

atoms in a metal.7.3.3 Explain the importance of alloys.

Guide for Reading

Build VocabularyWord Forms From what they know about the terms metallic and bond, have students infer the meaning of metallic bond.

Reading StrategyRelating Cause and Effect Have stu-dents describe the cause-and-effect relationship between metallic bonding and the properties of metals.

INSTRUCT

Have students read the text that opens the section. Ask, What property makes metals good electrical con-ductors? (Metal bonding involves highly mobile electrons, which are shared by all of the nuclei in a metallic solid. Electron mobility accounts for the high electrical conductivity of metals.)

Metallic Bonds and Metallic PropertiesRelateTo assess students’ prior knowledge about metals, ask, What are the prop-erties of malleable and ductile met-als? (They can be hammered into different shapes and drawn into wires.) What is an alloy? Give an example. (An alloy is a mixture of two or more ele-ments, at least one of which is a metal. Examples include steel and brass.)

1

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Answers to...Figure 7.12 repulsions between ions of like charge

Section Resources

Connecting to Your World

Section 7.3 Bonding in Metals 201

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Metalcation

Sea ofelectrons

Metal

Force

Strongrepulsions

Ionic crystal

Force

Metalrod

Die

Wire

Force

a b

7.3 Bonding in Metals

Guide for Reading

Key Concepts• How can you model the valence

electrons of metal atoms?• How are metal atoms arranged?• Why are alloys important?

Vocabularymetallic bonds

alloys

Reading StrategyUsing Prior Knowledge Before you read, jot down three things you know about metals. When you have read the section, explain how what you already knew helped you learn something new.Metallic Bonds and Metallic Properties

Metals are made up of closely packed cations rather than neutral atoms.The valence electrons of metal atoms can be modeled as a sea of elec-

trons. That is, the valence electrons are mobile and can drift freely fromone part of the metal to another. Metallic bonds consist of the attraction ofthe free-floating valence electrons for the positively charged metal ions.These bonds are the forces of attraction that hold metals together.

The sea-of-electrons model explains many physical properties of met-als. For example, metals are good conductors of electrical current becauseelectrons can flow freely in them. As electrons enter one end of a bar ofmetal, an equal number leave the other end. Metals are ductile—that is,they can be drawn into wires, as shown in Figure 7.12. Metals are also mal-leable, which means that they can be hammered or forced into shapes.

Figure 7.12 A metal rod can be forced through a narrow opening in a die to produce wire. As this occurs, the metal changes shape but remains in one piece. If an ionic crystal were forced through the die, it would shatter. Interpreting Diagrams What causes the ionic crystal to break apart?

a

b

You have probably seen deco-rative fences, railings, or weathervanes made of a metal called wrought iron. Wrought iron is a very pure form of iron that contains trace amounts

of carbon. It is a tough, malleable, ductile, and corrosion-resistant material that melts at a

very high temperature. As you already know, metals often have distinctive, use-ful properties. In this section, you will learn how metallic properties derive from

the way that metal ions form bonds with one another.

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Animation 9 See how metallic bonding explains some physical properties of metals.

The copyright holder has not granted permission todisplay this image in electronic format. Please seethe teacher's edition of your textbook for this image.

202 Chapter 7



Metals vs. Ionic CompoundsPurpose Students compare copper metal and a copper compound.

Materials copper metal or alloy, copper-containing ionic compound

Procedure Show the class a small sam-ple of elemental copper or a copper alloy and a sample of a copper-containing crystalline ionic mineral such as chalcocite (Cu2S). Wearing safety glasses and standing far from the students, smash both samples with a hammer. Discuss why the two sub-stances respond differently to the stress of the hammer blow.

Expected Outcomes The elemental copper will flatten but not break because the cations and electrons are mobile; the crystal will shatter.

Crystalline Structure of Metals

Use VisualsFigure 7.14 Lead a class discussion on the concept of “closest packing” of metal cations in pure metals. Use the three different closest packing arrangements shown in Figure 7.14 as a reference. Emphasize that the con-cept of closest packing also relates to more than just metal atoms. Have the students describe other examples of closest packing.

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Gifted and TalentedHave students develop formulas for calculat-ing the density of a metal given the atomic radius of the metal and the cubic unit cell packing arrangement of the metal atoms. Use sodium, potassium, iron, chromium, and

tungsten for body-centered cubic metals. Use copper, silver, gold, aluminum, and lead for face-centered cubic metals. The students may need to make models to see geometric relationships more clearly.

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202 Chapter 7

Hexagonal close-packed

ZincGold

Face-centered cubic

Chromium

Body-centered cubic

a b c

Both the ductility and malleability of metals can be explained in termsof the mobility of valence electrons. A sea of drifting valence electrons insu-lates the metal cations from one another. When a metal is subjected topressure, the metal cations easily slide past one another like ball bearingsimmersed in oil. In contrast, if an ionic crystal is struck with a hammer, theblow tends to push ions of like charge into contact. They repel, and thecrystal shatters.

Crystalline Structure of MetalsThe next time you visit a grocery store, take a look at how the apples ororanges are stacked. More than likely, they will have a close-packedarrangement, as shown in Figure 7.13. This arrangement helps save spacewhile allowing as many oranges to be stacked as possible.

Similar arrangements can be found in the crystalline structures of met-als. You may be surprised to learn that metals are crystalline. In fact, metalsthat contain just one kind of atom are among the simplest forms of all crys-talline solids. Metal atoms are arranged in very compact and orderlypatterns. For spheres of identical size, such as metal atoms, there are sev-eral closely packed arrangements that are possible. Figure 7.14 shows threesuch arrangements: body-centered cubic, face-centered cubic, and hexag-onal close-packed arrangements.

In a body-centered cubic structure, every atom (except those on thesurface) has eight neighbors. The metallic elements sodium, potassium,iron, chromium, and tungsten crystallize in a body-centered cubic pattern.In a face-centered cubic arrangement, every atom has twelve neighbors.Among the metals that form a face-centered cubic lattice are copper, silver,gold, aluminum, and lead. In a hexagonal close-packed arrangement, everyatom also has twelve neighbors. Because of its hexagonal shape, however,the pattern is different from the face-centered cubic arrangement. Metalsthat have the hexagonal close-packed crystal structure include magne-sium, zinc, and cadmium.

Checkpoint What metals crystallize in a face-centered cubic pattern?

Figure 7.13 These oranges illustrate a pattern called a hexagonal close-packed arrangement.

Figure 7.14 Metal atoms crystallize in characteristic patterns. Chromium atoms have a body-centered cubic arrangement. Gold atoms have a face-centered cubic arrangement. Zinc atoms have a hexagonal close-packed arrangement. Inferring Whichof these arrangements is the most closely packed?

a

b

c

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Alloys


Types of AlloysPurpose Students compare models of interstitial and substitutional alloys.

Materials styrofoam balls of different sizes and colors, toothpicks

Procedure Use styrofoam balls to illus-trate the crystal structures of interstitial and substitutional alloys. Use tooth-picks to hold the “atoms” together. Point out that brass is a substitutional alloy in which copper atoms are replaced by similarly sized zinc atoms. Steel is an interstitial alloy in which relatively small carbon atoms occupy the interstices between closely packed iron atoms.

Expected Outcome Students should be able to distinguish between the two types of alloys.

ASSESSEvaluate UnderstandingAsk, How can the conductivity of metals be explained? (Electrons can flow in and out of a metal because the valence electrons are not fixed.)

ReteachCompare chemical bonding in ionic compounds and pure metals.

Metal cations are surrounded by free-floating electrons. When metals are hammered, the cations move past each other. Conductivity results from mobile electrons.

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If your class subscribes to the Interactive Textbook, use it to review key concepts in Section 7.3.

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Answers to...Figure 7.14 face-centered cubic

Checkpoint

copper, silver, gold, aluminum, and lead

Section 7.3 Assessment23. metal cations surrounded by a sea of

mobile valence electrons24. Atoms in metals are arranged in a com-

pact and orderly manner.25. The properties of alloys are often supe-

rior to their component elements.26. ductile: can be drawn into wires; mallea-

ble: can be hammered into different shapes

27. Under pressure, the cations in a metal slide past each other. The ions in ionic crystals are forced into each other by the rigid structure.

28. body-centered cubic; face-centered cubic; hexagonal close-packed

29. Sterling silver used in jewelry is 92.5% sil-ver and 7.5% copper; bronze used in casting is 7 parts copper and 1 part tin.

Section 7.3 Bonding in Metals 203

Table 7.3

Composition of Some Common Alloys

Name

Composition

(by mass)

Sterlingsilver

Ag 92.5%Cu 7.5%

Cast iron Fe 96%C 4%

Stainlesssteel

Fe 80.6%Cr 18.0%C 0.4%Ni 1.0%

Springsteel

Fe 98.6%Cr 1.0%C 0.4%

Surgicalsteel

Fe 67%Cr 18%Ni 12%Mo 3%

AlloysAlthough every day you use metallic items, such as spoons, very few ofthese objects are pure metals. Instead, most of the metals you encounterare alloys. Alloys are mixtures composed of two or more elements, at leastone of which is a metal. Brass, for example, is an alloy of copper and zinc.

Alloys are important because their properties are often superior to thoseof their component elements. Sterling silver (92.5% silver and 7.5% copper)is harder and more durable than pure silver but still soft enough to be madeinto jewelry and tableware. Bronze is an alloy generally containing sevenparts of copper to one part of tin. Bronze is harder than copper and moreeasily cast. Nonferrous (non-iron) alloys, such as bronze, copper-nickel,and aluminum alloys, are commonly used to make coins.

The most important alloys today are steels. The principal elements inmost steel, in addition to iron and carbon, are boron, chromium, manga-nese, molybdenum, nickel, tungsten, and vanadium. Steels have a widerange of useful properties, such as corrosion resistance, ductility, hardness,and toughness. Table 7.3 lists the composition of some common alloys.

Alloys can form from their component atoms in different ways. If theatoms of the components in an alloy are about the same size, they canreplace each other in the crystal. This type of alloy is called a substitutionalalloy. If the atomic sizes are quite different, the smaller atoms can fit into theinterstices (spaces) between the larger atoms. Such an alloy is called an inter-stitial alloy. In the various types of steel, for example, carbon atoms occupythe spaces between the iron atoms. Thus, steels are interstitial alloys.


23. Key Concept How do chemists model the valence electrons in metal atoms?

24. Key Concept How can you describe the arrangement of atoms in metals?

25. Key Concept Why are alloys more useful than pure metals?

26. Describe what is meant by ductile and malleable.

27. Why is it possible to bend metals but not ionic crystals?

28. What are three different packing arrangements found in metallic crystals?

29. Describe two widely used alloys.

Figure 7.15 Bicycle frames are often made of titanium alloys that contain aluminum and vanadium.

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Explanatory Paragraph Write a paragraph describing how the sea-of-electrons model is used to explain the physical properties of metals.Hint: First write a sentence that summarizes the model. Then discuss how the model applies to specific properties of metals.

204 Chapter 7

Building With AlloysExplain that materials for a building must be able to withstand the stresses that a building undergoes. Discuss with students several different types of stress. Tensile stress can be observed when a building beam sags. The bot-tom of the beam undergoes tensile stress as it is slightly stretched. The top of the beam undergoes compression stress, which results when two forces push toward each other through a solid. The top of the beam is com-pressed as it is slightly shortened. A building undergoes shear stress in a strong wind. In shear stress, forces are applied from different directions, and the building might twist or break. Ask, Which of these buildings are most likely to undergo shear stress? (The Atomium, the Chrysler Building)

Metals are used as building materials because of strength and durability, but other properties might determine the use of a metal in a building. Ask, What property of a metal might make it useful on the outside of a building? (Answers might include luster when a shiny appearance is desired, malleability when the metal covers another material, or the ability of a certain metal to form a compound that protects the rest of the metal or other materials under it.) What metal was chosen for each these buildings? Why do you think that alloy was used? (possible answers: the Atomium—aluminum alloy; shiny, cor-rosion resistant, light, malleable; the Chrysler Building—steel; shiny, mallea-ble, corrosion-resistant; the Jewish Museum Berlin—zinc-titanium alloy: corrosion resistant, light)

A Good FoundationConcrete itself is not strong enough to be a good framing material. Reinforcing the con-crete by pouring it over steel rods that have been laid out in a grid adds a considerable amount of strength to the concrete. When exceptionally strong concrete is needed, it is

poured over steel cables that are stretched. After the concrete dries, the rods are released, and the concrete is compressed as the rods return to their original length. This exceptionally strong concrete is called pre-stressed concrete.

Facts and Figures

204 Chapter 7

Building with Alloys

Modern architecture would be a lot shorter if it weren’t

for steel. Since the late 1800s, using steel columns and

girders in construction has allowed architects to design

taller, stronger, and lighter buildings. Unlike buildings

made of wood, brick, or stone, steel structures are

strong enough to accommodate large, open interior

spaces that do not require supporting walls. Usually,

you can’t see the steel used to construct a building;

either it’s hidden by the floors and walls, or—in the case

of a building made of reinforced concrete—it’s actually

embedded in the floors and walls. The exteriors of

buildings often feature lighter alloys, such as alloys

of aluminum or titanium that resist corrosion.

Comparing and Contrasting How does

steel-frame construction differ from

reinforced-concrete construction?

The Atomium Brussels, BelgiumDesigned to resemble a crystal of iron magnified 165 billion times, the Atomium consists of nine spheres made of aluminum-alloy panels and connected by steel tubes. The top sphere contains an observation deck 92 meters above ground.

Jewish Museum BerlinBerlin, GermanyThis angular building is covered in thin sheets of zinc-titanium alloy. The untreated alloy will slowly oxidize and change color from exposure to the air and weather.

Chrysler BuildingNew York CityCompleted in 1930, this steel-frame high-rise stands 319 m tall and features a distinctive spire sheathed in shiny stainless steel.



Making an AlloyPurpose Students observe how to make an alloy from copper and zinc.

Materials penny, fine sandpaper, gran-ulated zinc, dilute NaOH solution, evaporating dish, tongs, hot plate, teaspoon

Safety Be sure to use adequate ventila-tion and have no skin contact with the NaOH solution. Do not touch hot objects.

Procedure Use the sandpaper to clean any tarnish from the penny. Add a tea-spoon of zinc to the evaporating dish, and cover the zinc with NaOH solution. Place the penny on the zinc, being sure the penny is also covered by the solu-tion. Heat the dish until the penny changes to a silvery color. Using the tongs, remove and rinse the penny. Place the penny on the hot plate, which should be set to medium heat. When the penny turns a gold color, use tongs to remove it from the hot plate.

Expected Outcome A gold-colored alloy forms from copper and zinc.Ask, Why did the penny turn a silver color? (Zinc was deposited on the penny.) At what point was an alloy formed? (An alloy was formed when the copper and zinc on the penny were heated to a golden color.) Why was the penny heated to form the alloy? (Heat increases the kinetic energy of the atoms, allowing them to mix more freely.)

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Answers to...Comparing and Contrasting Steel frames allow for more flexible inte-rior spaces. Reinforced-concrete construction takes up less vertical space, so more floors can be con-tained in a building of a certain height if reinforced concrete is used instead of steel.

Less Proficient ReadersHave students create a list of the building materials they find in their school building. If the material is a metal, have them try to determine whether the metal is an alloy or

not. Point out to students that most of the metals will be alloys, such as steel, because alloys can be made to have the properties desired in a building material.

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Technology and Society 205

Reinforced-concrete constructionAnother method for building high-risesis to use reinforced concrete. Concrete ismore rigid than steeel and holds up betterstructurally in fires. In addition, a concreteframe takes up less vertical space than asteel frame, meaning that a concrete-framebuilding can contain more floors than asteel-frame building of the same height.

Steel-frame constructionThe most common methodto build a high-rise is to usea steel framework. Steelframes can be assembledquickly and allow forflexible interior spaces.

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