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

Подождите немного. Документ загружается.

Chemical Applications and Practices

41. Sodium and potassium are often found as ions in sports

drinks. Which is the larger element? Explain.

42. Helium and hydrogen have both been used in dirigibles (also

known as blimps or airships). Which of these two element’s

atoms is larger? Explain.

Section 7.6 Ionization Energies

Skill Review

43. For this calculation, use the ﬁrst ionization values of sodium

and potassium found in the chapter. How many grams of

potassium could be ionized using the same energy that is re-

quired to ionize 1.00 g of sodium?

44. For this calculation, use the ﬁrst ionization values of calcium

and potassium found in the chapter. How many grams of cal-

cium could be ionized using the same energy that is required

to ionize 5.00 g of potassium?

45. Within each of these two lists, rank the elements listed from

lowest to highest ﬁrst ionization energy:

a. Al, Si, P b. Ne, Ar, Kr

Review your rankings. Would you make any changes if asked

to do the rankings by second ionization energy? If, so de-

scribe the reasons for any changes.

46. Within each of these two lists, rank the elements listed from

lowest to highest ﬁrst ionization energy:

a. Li, Be, B b. K, Rb, Cs

Review your rankings. Would you make any changes if asked

to do the rankings by second ionization energy? If, so de-

scribe the reasons for any changes.

47. Based on the following ionization energy (I) information,

predict which element(s) are most likely metals and which

are nonmetals. Also, suggest the most likely periodic table

group for each of these “unknown” elements. (Note that

these values, all of which are in units of kilojoules per mole,

may not reﬂect actual ionization energies of known elements.)

disputed. If element 118 does exist, to which group in the

periodic table does it belong? Would its ionization energy be

expected to be higher or lower than that of its next neighbor

above it in the periodic table? Explain your answer.

50. Element 119 could be discovered in the near future. To which

group in the periodic table would it belong? Would its ion-

ization energy be expected to be higher or lower than that of

its neighbor directly above it in the periodic table? Explain

your answer.

51. The ﬂow of potassium and sodium ions into and out of nerve

cells makes it possible for signals to be sent throughout our

bodies. The fact that you are reading right now is dependent

on this ﬂow of ions.

a. What are the charges on the potassium and sodium ions,

respectively?

b. Which is the larger of the two ions, sodium or potassium?

c. Which of the two ions, sodium or potassium, is easier to

produce from their neutral atoms?

52. Calcium ions also have great importance in cells and move

into and out of other channels in cell membranes. What

would be the charge on a calcium ion? Would this ion be

larger or smaller than a potassium ion? Explain.

53. Explain why the ﬁrst ionization energy of sodium is lower

than magnesium’s ﬁrst ionization energy, but the second ion-

ization energy of sodium is higher than that of magnesium.

54. Examining the ﬁrst ionization energies of the elements in

Period 4 of the periodic table reveals the general tendency for

an increase in ionization energy from left to right. However,

there is a slight drop in ionization energy from calcium to

gallium. Study the electron conﬁguration of both of these el-

ements and offer an explanation for this apparent deviation

from the trend.

55. Many elements may be found in nature in a variety of oxida-

tion states. However, these statements express some limita-

tions. Explain the basis for each of these limitations.

a. Magnesium is never found naturally as the +3 ion.

b. Fluorine is never found naturally as the +1 ion.

c. Hydrogen is never found naturally in the +2 oxidation

state.

d. Aluminum tends to “prefer” the +3 ion naturally.

56. Explain the basis for each of these limitations on the oxida-

tion states in which certain elements appear.

a. Arsenic can often be found with a charge of +5 but never

with a charge of +6.

b. Titanium could be found with either a +2 or a +4 charge

but not with a +5 charge.

c. Potassium is never found naturally as a neutral element.

d. Tin can be found with a +2 or a +4 charge but not with

a +5 charge.

Section 7.7 Electron Afﬁnity

Skill Review

57. Which of these would have the best chance of forming a

stable anion: S or Xe? Justify your choice.

58. Which of these would have the best chance of forming a

stable anion: Cl or Ar?

59. Which of these has the more exothermic value for electron

afﬁnity: Cl or Ar?

298 Chapter 7 Periodic Properties of the Elements

(kJ/mol) I

(kJ/mol)

Element A 500 4600 5800 7460

Element B 1060 1900 2900 5000

(kJ/mol) I

(kJ/mol)

Element A 740 1440 7730 8670

Element B 2100 3230 4400 5500

48. On the basis of the following ionization energy (I) informa-

tion, predict which element(s) are most likely to be metals

and which to be nonmetals. Also, suggest the most likely pe-

riodic table group for each of these “unknown” elements.

(Note that these values, all of which are in units of kilojoules

per mole, may not reﬂect actual ionization energies of known

elements.)

Chemical Applications and Practices

49. Some claims have been made regarding the production (dis-

covery) of element 118. However, those claims have been

60. Which of these has the more exothermic value for electron

afﬁnity: O or F?

61. Arrange the following atoms on the basis of electron afﬁnity,

from least negative to most negative: Na, Mg, Al. Now

arrange them on the basis of their ability to form anions. Is

the order any different? If so, explain any changes you made

in the ordering.

62. Arrange the following list in order from least negative elec-

tron afﬁnity value to most negative electron afﬁnity value:

Br, Br

−

,K.

63. Because they all have a negative charge, the electrons in

atoms repel each other. Explain why the electron repulsion

in a ﬂuorine atom is such a large factor in determining its

unusually small electron afﬁnity value.

64. Explain why ionization energy is always a positive quantity,

whereas electron afﬁnity may be positive or negative.

Chemical Applications and Practices

65. The ﬁrst ionization energy of sodium is +495 kJ/mol. The

electron afﬁnity of chlorine is −349 kJ/mol. When sodium

metal and chlorine gas are placed near each other, a violent

reaction takes place. After a sodium atom has lost an electron

and a chlorine atom has gained an electron, both are oppo-

sitely charged and have the stable conﬁguration of noble

gases.

a. What is the total energy change that occurs in these two

processes? Is this exothermic or endothermic?

b. Two additional energy changes occur in this reaction: the

dissociation of a chlorine atom from Cl

(+158.8 kJ/mol)

and the energy released when the ions are combined

(−1030 kJ/mol). What can you determine about the spon-

taneous, violent reaction between sodium and chlorine

when all of the energy changes in this process are consid-

ered?

66. The ﬁrst ionization energy of lithium is 5.392 eV. (1 eV =

96.485 kJ/mol). The electron afﬁnity of ﬂuorine is

−328 kJ/mol. When lithium metal and ﬂuorine gas are

placed near each other, a reaction takes place. After the ion-

ization of lithium and the formation of anionic ﬂuorine,

both are oppositely charged and have the stable conﬁgura-

tion of noble gases.

a. What is the total energy change that occurs in these two

processes? Is this exothermic or endothermic?

b. Two additional energy changes occur in this reaction: the

dissociation of a chlorine atom from F

(+158.8 kJ/mol),

and the energy released when the ions are combined

(−1030 kJ/mol). What can you determine about the

spontaneous, violent reaction between lithium and ﬂuo-

rine when all of the energy changes in this process are

considered?

Section 7.8 Electronegativity

Skill Review

67. Arrange this list of atoms in a correct ranking from least elec-

tronegativity to highest electronegativity: Na, F, As, Li, S.

68. Arrange this list of atoms in a correct ranking from least elec-

tronegativity to highest electronegativity: K, P, O, Br, N.

69. Each of these situations depicts two atoms bonded to each

other. In each case, which atom is more likely to attract elec-

trons toward itself within the bond?

a. NOOb.SeOSc.BrOGe d. ClOO

70. Each of these situations depicts two atoms bonded to each

other. In each case, which atom is more likely to attract elec-

trons toward itself within the bond?

a. FOOb.POSc.HOCd.COS

Chemical Applications and Practices

71. Using the information in the chapter, determine the change

in electronegativity from Ga to Se. This change arises as the

number of protons is increased by three. The change in

proton number from Sc to Zn is nine. What is the change in

electronegativity from Sc to Zn? Based on the change in

number of protons, is this what you would expect? Why or

why not?

72. In an early section we noted that the trend in electron afﬁn-

ity from ﬂuorine to chlorine was a bit different than we

might expect. The trend in electronegativity is, however,

what we would expect; that is, it decreases from top to

bottom in the group. What aspect of the deﬁnition of elec-

tronegativity helps explain the difference between the two

trends?

Section 7.9 Reactivity

Skill Review

73. Arrange this list of metals in order from generally least reac-

tive to most reactive: Na Mg Rb

74. Arrange this list of nonmetals in order from generally least

reactive to most reactive: S Cl I

75. What is the relationship between ﬁrst ionization energy and

reactivity in metals and in nonmetals?

76. What is the relationship between electron afﬁnity and reac-

tivity in metals and in nonmetals?

Chemical Applications and Practices

77. The “coinage metals” are copper, silver, and gold.

a. Describe the location of these elements in the periodic

table.

b. Where are they located in the activity series of metals

used in the chapter?

c. Is the location of Group IB

consistent with periodic

reactivity and the

activity series?

Focus Your Learning

299

Coins made from copper,

silver, and gold. In the past,

your pocket might contain

coins made from these

three metals. Due to the

value of silver and gold, the

U.S. mint currently uses

other metals and mixtures

of metals known as alloys.

78. Aluminum metal is considered a fairly active metal. It cer-

tainly reacts with oxygen more vigorously than does iron.

The reaction of both metals with oxygen produces oxides

that have different characteristics. Contrast the properties of

the two oxides.

79. Nonmetals such as chlorine and oxygen are both considered

very reactive. However, they may also react with each other.

Using any available information, decide which of the two

would probably become more negative in the reaction? Ex-

plain the basis for your answer.

80. Until the 1960s the noble gases were considered chemically

inert, or totally unreactive. Eventually, however, some noble

gases were made to react. Among the ﬁrst compounds of

noble gases, ﬂuorine was typically a reactant. Explain why

ﬂuorine was such a good candidate for this role of reactivity.

Section 7.10 The Elements and the Environment

Skill Review

81. What explanation can be given for the fact that iron is found

in much greater abundance in the Earth’s crust than in the

mantle?

82. What substance that is common in the Earth’s crust accounts

for most of the silicon and oxygen (the two highest-ranking

elements in the crust?)

83. a. What is the third most abundant substance, by volume, in

the Earth’s atmosphere?

b. How many particles of that substance would be found in

one mole of dry air?

84. a. What is the fourth most abundant substance, by volume,

in the Earth’s atmosphere?

b. How many molecules of that substance would be found

in one mole of dry air?

Chemical Applications and Practices

85. Table 7.15 lists the ranked abundance of elements in the

Earth’s crust by mass. Which element(s) in that list are most

likely to be found in an uncombined state? (Uncombined,in

this case, means “not combined with other elements.”)

86. In Table 7.15, titanium is listed as more abundant, by mass,

than hydrogen. Convert the mass to moles and number of

atoms. If the ranking were to be redone by number of atoms,

would hydrogen still be below titanium? Prove your answer.

Comprehensive Problems

87. For each of these, provide a brief summary of his contribu-

tion to organizing the properties of elements into a useful

arrangement.

a. Johan Dobereiner

b. John Newlands

c. Dmitri Mendeleev

88. Indium (element 49) is an important element, but it is

not commonly used in textbook examples. Use the Internet

or another text reference to determine a chemical and a

physical property that indium has in common with gallium,

element 31.

89. a. What is the most common form of steel used in the

world?

b. Based on composition, what is the chief difference

between this most common type of steel and other types

of steel?

90. Based on the characteristics listed, which type of element

(metal, nonmetal, or metalloid) is being described?

a. At room temperature the sample is a dull, brittle solid.

The element is most likely to assume a negative charge

in ionic compounds.

b. At room temperature the sample is a solid that is able to

conduct electricity.

91. Given the information depicted in the illustration below,

what would you predict for the bond length between a

carbon and a hydrogen atom?

92. Contrast the meaning of the term valence electrons in the

context of calcium and in the context of chromium.

93. a. Using the ﬁrst 18 elements on the periodic table, prepare

a graph with number of valence electrons on the y axis

and group number on the x axis. Is this a periodic func-

tion? Explain why or why not.

b. Prepare another graph of the same elements, using the

most common oxidation number on the y axis and the

group number on the x axis. Is this a periodic function?

Explain why or why not.

94. The deﬁnition of the atomic radius of an atom is basically

straightforward; it is one-half the distance between the

nuclei of a molecule made of two identical atoms.

a. What problems would arise if we were to deﬁne the

radius as one-half the diameter of an atom?

b. How does metallic radius differ from covalent radius?

95. Explain why, when comparing the radius of two atoms, it

would be important not to base the comparison solely on

which atom has the greater number of protons.

96. a. This reaction depicts an atom being ionized. To which

side of the reaction (left or right) should you show the

energy term for the reaction?

X → X

+ e

−

b. Select the correct response for each of these situations:

The smaller the value for ionization energy, the (less,

more) easily ionization will occur. The smaller the value

for ionization energy, the (less, more) chemically reactive

a metal will be.

97. Seldom do we get to describe a chemistry situation as

“always” happening without exception. However, this state-

ment is always true: Successive ionization energies in an

atom are always higher than previous values. What underly-

ing factors make this statement true?

98. Electron shielding or electron screening provides a descriptive

way to explain how the attraction for an electron by the nu-

cleus can vary in different electron arrangements. However,

H–H bond

75 pm

C–C bond

154 pm

300 Chapter 7 Periodic Properties of the Elements

electron screening can have a slightly different description

when applied to elements in the same row on the periodic

table than when applied to elements in the same group.

Contrast the type of screening, and its effectiveness, when

the term is applied to a period and when it is applied to a

group.

99. Arsenic and selenium are next to each other in Period 4 of

the periodic table. In general the trend, from left to right on

the table, is in favor of an increase in ionization energy.

However, when you compare the values for As and Se, you

note that the ionization drops slightly instead of increasing.

What is the basis for this situation?

100. Radioactive fallout from nuclear testing can contain signif-

icant amounts of unstable strontium. This can be particu-

larly harmful to young children, who have rapidly growing

bone structures. What common charge would strontium

ions have? What element is an important component in

bone development? Why is exposure to radioactive stron-

tium such a grave danger to young children?

101. a. The electron afﬁnity of chlorine is approximately

−350 kJ/mol. Does this indicate an exothermic or an en-

dothermic reaction?

b. Write out the reaction, including placing the energy term

in the equation, that depicts the “electron afﬁnity reac-

tion” for the process described in part a.

102. When summarizing the periodic trends in ionization

energy, atomic radius, electronegativity, and (to some

extent) electron afﬁnity, we can typically indicate the trend

with a one-directional arrow to show the trend increasing

or decreasing from left to right in a row, or from top to

bottom in a group. Explain why the trends for increasing

reactivity start from the center of the periodic table and

move outward in both directions.

103. In the chapter, we note that tool steel contains 4.38% by

mass of chromium and 0.864% by mass of carbon.

a. What is the electron conﬁguration of chromium?

b. If a steelmaker wishes to make 1.0 kg of tool steel, which

element (carbon or chromium) would require a larger

number of moles?

c. How many atoms of chromium would be present in an

88.7 g sample of tool steel?

104. The halogens are so-named because they react with metals

to make salts.

a. How many valence electrons do each of the atoms in this

group contain?

b. Write the balanced reaction that occurs between sodium

metal and iodine crystals.

105. According to Table 7.13, the human body contains the same

mass percent of silicon and potassium. Which occurs in the

body in a greater number of moles?

106. If the periodic table was arranged, in order, from smallest

atom (based on atomic radius) to largest atom

a. What atom would be listed ﬁrst?

b. What atom would be listed last?

c. Would the noble gases still be aligned in a column?

d. How many helium atoms would be needed to create a

line of atoms 1.0 inches long (assuming that the atoms

are just touching and placed in a straight line)?

107. In the chapter, we mention that francium-210 can be made

from gold-197 by bombarding the gold with oxygen-18.

a. How many protons, neutrons, and electrons are found in

one atom of francium-210?

b. How many grams of francium-210 could be made from

2.50 g gold-197?

c. Which has a larger atomic radius—francium-210 or

gold-197? (Assume speciﬁc isotopes do not differ in their

atomic radius.)

Thinking Beyond the Calculation

108. An aqueous solution of sodium bicarbonate (NaHCO

)

reacts with aqueous hydrochloric acid to produce carbonic

acid (H

) and sodium chloride (NaCl).

a. Write a balanced equation illustrating this reaction.

b. Indicate the group and period for each of the elements

involved in the reaction.

c. Indicate those elements in the reaction that can be char-

acterized as metals.

d. Carbonic acid decomposes in solution to produce carbon

dioxide and water. Write the balanced equation for this

reaction.

e. If 10.0 mg of sodium bicarbonate is added to 275 mL of

water and reacted completely with a stoichiometric

amount of HCl, what concentration (in ppm) of sodium

chloride will result? (Assume no volume change.)

f. Using the information from part e, determine what the

concentration (in ppm) of sodium ions will be.

g. A pastry chef may wish to perform a reaction similar to

this using sodium bicarbonate and acid. Why would the

chef wish to perform this reaction?

Focus Your Learning

301

Reaction of sodium bicarbonate and HCl.

Bonding

Basics

Bottles in the medicine chest. The active

ingredients in these medications are a

mix of molecules and ionic compounds,

two classes of compounds that differ in

the way their atoms are bound together.

302

Contents and Selected Applications

8.1 Modeling Bonds

8.2 Ionic Bonding

Chemical Encounters: The Uses and Behavior of Sodium Chloride

Chemical Encounters: Focus on Zeolites

Chemical Encounters: Fluoridated Water and Tooth Decay

8.3 Covalent Bonding

Chemical Encounters: Tetraethyl Lead in Gasoline

Chemical Encounters: Calcium Channel Blockers

8.4 VSEPR—A Better Model

Chemical Encounters: Focus on Morphine

8.5 Properties of Ionic and Molecular Compounds

Go to college.hmco.com/pic/kelterMEE for online learning resources.

Americans spend more than $250 billion per

year on prescription and over-the-counter (OTC)

medicines, about one-half of the world’s total. Our

medicine cabinets are fairly well stocked. If you and your

family are typical consumers, you may have headache pills,

muscle relaxers, antacids, cough syrup, and a couple of old bottles

of antibiotics. A quick check of the active ingredients reveals that these

bottles contain such compounds as ibuprofen, magnesium salicylate, sodium

bicarbonate, and pseudoephedrin. One of the more common ingredients in a

pain-relieving OTC medicine is aspirin (a molecular compound). Sodium bicarbon-

ate (an ionic compound) is typically used as an antacid. These two compounds

work in different ways to relieve common maladies, and they are fundamentally

different in their chemical makeup. They do not have the same numbers and types

of atoms, although this alone isn’t enough to explain why these compounds differ

radically in many properties, including melting point, boiling point, solubility in

water, and chemical reactivity. The main difference lies in their designation as mole-

cular compounds or ionic compounds, and this designation is made on the basis of

the way in which the atoms are bound together.

Biochemists, medicinal chemists, and

pharmacognocists (who study the properties

of drugs, especially focusing on those from natural sources) are familiar with molec-

ular and ionic compounds. One of the fundamental questions asked by these scien-

tists is

How do the bonds within a compound, the shape of the compound, and the physical

properties of the compound contribute to its ability to treat a disease or common malady?

In this chapter, we will answer this question and others as we explore how atoms

are held together, and how this bonding determines the shapes and properties of

compounds.

8.1 Modeling Bonds

We will begin our discussion by looking at the ways in which researchers

view the compounds with which they work. Although the formula of a po-

tential drug can be useful for determining the number and types of atoms it

contains, it doesn’t say much about the molecular shape. Aspirin—acetylsal-

icylic acid—which we discussed in Chapter 3, has the formula C

.The

formula alone provides no information about how aspirin structurally inter-

acts with the body to alleviate a headache. To ﬁnd out more, we must look

deeper into the atom—into the role that the electron plays in determining

shapes. Our ﬁrst key idea is that the different arrangements of the electrons in a

compound help determine the shape and the properties of that compound.

Three Kinds of Bonds

In their search for marketable pharmaceuticals, biochemists, medicinal

chemists, and pharmacognocists use a wide range of

molecular models, three-

dimensional depictions of the structure of molecules, as tools to examine the

shape and properties of compounds. To be ideal for the marketplace, a com-

pound should

■

effectively treat a particular disease, malady, or ailment

■

have no serious side effects

■

be inexpensive to mass-produce

Acetylsalicylic acid

303

The molecular models that researchers use in their search for marketable drugs

serve as representations of the

chemical bonds, the forces that hold atoms together,

within a compound. Chemical bonds arise between atoms when some of the out-

ermost electrons on the bonding atoms interact. In some cases, the electrons tend

to congregate on one of the atoms of the bond. In other cases, the electrons are

shared more or less equally between the atoms. A vital point to keep in mind is

that there are no absolutes. There are many degrees of electron sharing, from more

or less complete ownership by one atom to about equal sharing by both atoms,

and every possibility on the scale of sharing can occur. Still, we classically think

of three types of chemical bonds: the covalent bond, the ionic bond, and the

metallic bond.

Sodium chloride, the stuff we sprinkle on our French-fries, is an example of a

compound with an ionic bond. The

ionic bond lies at one end of the spectrum of

chemical bonds, where the atoms are held together by the force of attraction of

opposite charges. We say that one or more electrons have been removed from one

atom (remember, there are no absolutes!) and congregate on the other atom in

the compound.

Aspirin is an example of a compound in which the atoms are held together by

covalent bonds. The

covalent bond lies at the other end of the spectrum of chem-

ical bonds, where electrons are shared between the atoms. The positively charged

nuclei on either end of the bond attract the negatively charged electrons. It is this

force that holds the atoms together.

The aluminum atoms in a soda can are held together with a metallic bond.

The

metallic bond is a special type of bond in which metal cations are spaced

throughout a sea of mobile electrons. We’ll learn more about this bonding

pattern in Chapter 13.

Lewis Dot Symbols

The ﬁrst step in the construction of the molecular model of a compound, such as

aspirin, sodium chloride, or aluminum, is drawing an atom itself. We recognize

that the “business part” of the atom is its set of valence electrons. In 1916,

G. N. Lewis (Figure 8.1) developed a useful shorthand representation employing

Covalent and ionic bonds lie at opposite ends of the bonding spectrum. Real

bonds between different elements are neither purely ionic nor purely covalent.

Purely

covalent

Purely

ionic

“Real” bonds

304 Chapter 8 Bonding Basics

FIGURE 8.1

The American chemist G. N. Lewis

(1875–1946) developed a model of

atomic bonding that is still used today.

These notes, written by Lewis in 1902,

illustrate his thinking on how electrons

were arranged around an atom. The

Lewis dot symbols that we use today are

slightly modiﬁed from this original work.

Visualization: Covalent Bonding

Tutorial: Covalent and Ionic

Bonding

Lewis dot symbols. Typically, the ﬁrst two dots, representing valence electrons in

the s orbital, are placed one at a time, as a pair, to the side of the element symbol.

The next three dots, representing valence electrons in the p orbitals, are placed

individually on the other sides of the element symbol. The remaining electrons

are placed alongside of the ﬁrst three dots to indicate pairs of electrons in the

p orbitals. The Lewis dot symbols for the elements of Period 2 are shown below.

Lewis dot symbols can be used to represent ions as well. Because an ion is just

an atom (or a group of atoms) with a different number of electrons than the un-

charged species, we can draw its Lewis dot symbol. For instance, the ﬂuoride ion

is a ﬂuorine atom plus an electron (F

−

). To distinguish ions from Lewis dot sym-

bols of atoms, we place brackets around the drawing, and we place the charge on

the ion outside the brackets. Note in the diagram below that the nitrogen anion

has three extra electrons and the sulfur anion has two extra electrons. All three

have eight valence electrons and the same electron conﬁguration as the period’s

noble gas (Ar for S and Ne for F and N). This idea is important in

Lewis dot

structures

, and we’ll revisit it later in this chapter.

Lewis dot symbols can also be drawn for cations. In a monatomic cation, elec-

trons have been removed from the atom (or group of atoms,) resulting in a pos-

itive charge. When we write the Lewis dot symbols, we indicate the number of

valence electrons and the resulting positive charge.

Electron Conﬁguration of Ions

We determined the electron conﬁguration of the elements in Chapter 6 using the

Aufbau principle. Electron conﬁgurations can also be used to describe which

atomic orbitals in ions contain electrons. For example, the electron conﬁguration

of the sodium ion contains one electron less than that of the sodium atom. This

ion, which is necessary for proper contraction of heart tissue and electrolyte

balance inside and outside of the body’s cells, has an electron conﬁguration that

lacks the valence electron from the sodium atom.

Na: 1s

:1s

In anions such as the ﬂuoride ion (F

−

, found in toothpaste and added to most

U.S. municipal water supplies to help prevent tooth decay), the electron conﬁgu-

ration shows the addition of another electron to the valence shell:

F: 1s

−

:1s

We also know from our discussion of ionization energy that valence electrons are

the most accessible of the electrons in an atom. The addition of an electron to an

atom to make an anion occurs in the valence shell. Similarly, cations can be made

by removing valence electrons.

How do the electron conﬁgurations of Na

and F

−

compare with each other and with the noble gas nearest to them in atomic number

on the periodic table?

They have the same electron conﬁguration as neon,

. We say that the electron conﬁguration of Na

is isoelectronic with (has

the same electron conﬁguration as) F

−

and Ne.



3

N F



32

NeF

ONCBBeLi

8.1 Modeling Bonds 305

Lewis dot symbols of elements 3 to 10.

Lewis dot symbols of some

monatomic anions.

Lewis dot symbols of cations.

Video Lesson: Valence Electrons

and Chemical Bonding

The position of the electrons as either valence or core electrons on an atom

can help us understand the structure of a molecule such as aspirin. Knowing

where the electrons on the atom reside and how those electrons behave is of ut-

most importance in developing our three-dimensional model of a molecule. As

we’ll see shortly, the number of valence electrons on the atom is also important to

building our best possible model.

Octet Rule

In any chemical reaction, one of the driving forces is the ability of each atom to

reach a stable electron conﬁguration. Electrons shufﬂe around until each atom

has its most energetically stable arrangement of electrons. As in F

−

and Na

, the

most stable electron conﬁguration of the main-group elements is isoelectronic

with a noble gas. We refer to this as the

octet rule because the stable arrangement

for all of the noble gases beyond helium has eight valence electrons. For the ele-

ments H and He and the ions Li

and Be

, the rule is also known as the duet rule

because of the need for only two electrons to ﬁll the valence shell of the ﬁrst row

elements. In other words, main-group atoms typically react by changing their num-

ber of electrons in such a way as to acquire the more stable electron conﬁguration of

a noble gas. A full valence shell around an atom is a good situation for an atom,

because it then has the same electron conﬁguration as a noble gas. We’ll use this

rule a lot as we put together atoms to make ionic compounds and molecules.

EXERCISE 8.1 Writing Lewis Dot Symbols

Write the electron conﬁguration, and the Lewis dot symbol, for both P

3−

and Al

With which element are they isoelectronic?

Solution

The P

3−

anion has three electrons more than the phosphorus atom. Therefore, the

electron conﬁguration is 1s

. This is the same electron conﬁguration as

argon. The Lewis dot symbol shows an octet of electrons.

The Al

cation is missing three electrons, compared to the neutral atom. Its

electron conﬁguration is 1s

, which is isoelectronic with neon. The Lewis dot

symbol also illustrates an octet of electrons.

PRACTICE 8.1

Write the electron conﬁguration, write the Lewis dot symbol, and determine which

atoms are isoelectronic with S

2−

−

,Mg

, and Br

−

See Problems 1–4 and 9–14.

8.2 Ionic Bonding

We noted before that sodium chloride is an ionic compound. In ancient times, it

was a highly sought-after seasoning used in cooking and pickling. According to

the United States Geological Survey (USGS), 210 million metric tons of sodium

chloride were harvested from salt water or mined from deposits in the ground in

2005 (see Figure 8.2). In the United States, most of the salt that is harvested is

used to manufacture chlorine gas and sodium hydroxide. About 37% of the total

is used to de-ice highways. Only 3% is used in the food industry. Even so, the

average American consumes more than 1.5 kg of salt each year!

The human body requires only about 500 mg of sodium per day, yet the aver-

age American ingests between 2300 and 6900 mg each day. This high level of

306 Chapter 8 Bonding Basics

Application

HEMICAL

ENCOUNTERS:

The Uses and

Behavior of

Sodium Chloride

8.2 Ionic Bonding 307

sodium consumption can contribute to hypertension,

sleep apnea, and other disorders.

Sodium chloride (NaCl) is a compound held to-

gether by ionic bonds. As we saw in Chapter 4, sodium

chloride, like all other Group IA salts, dissociates

essentially completely into its ions when added to

water. It is a strong electrolyte and is a good example of

a typical ionic compound. Other examples of impor-

tant ionic compounds are shown in Table 8.1 and Fig-

ure 8.3. Note that some ionic compounds contain ionic

and covalent bonds. For example, calcium carbonate

contains an ionic bond between the calcium ion and

the carbonate ion. The carbon and oxygen atoms in the

carbonate ion are covalently bonded to each other. What makes sodium chloride

a “typical” ionic compound? The answer lies in the nature of its bonding and

properties, which we now explore.

FIGURE 8.2

Harvested sodium chloride is stored in

piles near the puriﬁcation plant.

Important Ionic Compounds

Compound Formula Selected Use

Calcium carbonate CaCO

Limestone, chalk

Calcium chloride CaCl

Sidewalk salt

Iron(III) oxide Fe

Pigment

Magnesium hydroxide Mg(OH)

Milk of magnesia

Sodium carbonate Na

Glass, soaps, and detergents

Sodium bicarbonate NaHCO

Baking soda

Sodium chloride NaCl Production of chlorine and sodium

hydroxide, seasoning, saline

solutions

Sodium ﬂuoride NaF Toothpaste

TABLE 8.1

FIGURE 8.3

Important ionic compounds.