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

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Description of Ionic Bonding

Sodium chloride is commercially mined or processed from brine (salt water, such

as from the oceans or the Great Salt Lake). In the laboratory, we can combine

sodium metal and chlorine gas in a violent reaction to make the salt. The reac-

tion, illustrated below, involves the transfer of electrons from sodium atoms to

Na(s) +

⁄

2Cl

(g) n NaCl(s) ∆H =−410 kJ

chlorine. If we look at this reaction more closely, we note that one of the atoms

(the metal) loses electrons to become a cation. In our reaction, the sodium atom

Na• n [Na]



+ e

−

loses an electron to become the sodium cation (Na

). The other atom (the non-

metal) gains electrons to become an anion. Addition of an electron to a chlorine

atom forms the chloride anion (Cl

−

). The charges on the sodium and chloride

ions attract each other, causing the ions to associate with each other in an

ion pair

(Na

−

But it doesn’t stop there. Other sodium cations are also attracted to the nega-

tive charge on the chloride. Other chloride anions are attracted to the positively

charged sodium. The end result is a collection of alternating sodium and chloride

ions arranged in a solid

crystalline lattice, a highly ordered, three-dimensional

arrangement of atoms, ions, or molecules (explained in detail in Chapter 13).

Within a sodium chloride crystal, each of the sodium cations has six neighboring

chloride anions, and each of the chloride anions has six neighboring sodium

cations. The forces of attraction combine to provide a neatly packed crystal of

alternating sodium and chloride ions, as shown in Figure 8.4.

Examples of Ionic Bonding

We can illustrate the formation of an ionic compound such as sodium chloride

(NaCl) through the use of Lewis dot symbols. As reﬂected in their positions on

the periodic table, sodium is a metal and chlorine is a nonmetal. Sodium has a

relatively low ionization energy and loses an electron to achieve the electron

conﬁguration of a noble gas.

Na• n [Na]



+ e

−

n 1s

Chlorine, which has a very favorable electron afﬁnity, gains an electron to achieve

a noble gas electron conﬁguration.

n 1s

In the formation of NaCl from Na metal and Cl

gas, the electron from the

sodium atom is added to the valence shell of the chlorine atom. We can represent

this process in equation form by showing the movement of the electron from the

sodium atom to the chlorine atom with a ﬁshhook-shaped arrow. The result is a

sodium cation and a chloride anion that are attracted to each other. The ionic

crystal that results has the formula NaCl.

Energy must be added to the system to remove the electron from the sodium

atom. Energy, however, is released when that electron is added to the chlorine atom,

and even more energy is released when the ionic bond forms. The net result from the



ClCl e







ClCl e





308 Chapter 8 Bonding Basics

–

FIGURE 8.4

Crystal structure of sodium chloride. In

the crystal structure of NaCl, we note

that each sodium ion is surrounded by

six chloride ions and that each chloride

ion is surrounded by six sodium ions. The

resulting crystal possesses a regular

close-packed pattern of ions (see

Chapter 14).

Video Lesson: Ionic Bonds

Visualization: Structure of an

Ionic Solid (NaCl)

8.2 Ionic Bonding 309

formation of a crystalline lattice of sodium and chloride ions is a large release of

energy.

EXERCISE 8.2 Ionic Compound Formation

Calcium chloride, CaCl

, is spread on our sidewalks to melt ice. Because of its abil-

ity to absorb moisture, it is used to dry gases and prevent dust buildup in mining

and highway maintenance. Using Lewis dot symbols, show the transfer of electrons

to form CaCl

from calcium metal and chlorine atoms.

Solution

We start by drawing the Lewis dot symbols for the calcium and chlorine atoms.

Then we show the transfer of an electron from the less electronegative calcium to

the more electronegative chlorine atom (arrow).We show a similar transfer with the

other calcium electron to a different chlorine atom. We end with three ions (Ca

and two Cl

−

) that have noble gas electron conﬁgurations. Note that the ions end

up with either a completely full valence electron conﬁguration (and become iso-

electronic with a noble gas) or a completely empty valence shell (and become

isoelectronic with a noble gas). This is the octet rule in action.

PRACTICE 8.2

Use the Lewis dot structure model to show the formation of sodium oxide from

atomic sodium and oxygen atoms.

See Problems 21, 22, and 27–30.

Medicinal chemists, biochemists, and pharmacognocists know about the dri-

ving force for the formation of ionic bonds.Ethidiumbromide,shown in Figure8.5,

is a toxic compound that is used by researchers to stain DNA in order to make it

easier to see. This compound exhibits its effects by inserting itself into DNA

strands, causing structural changes to the DNA molecule. The negative charges

on the DNA and the positive charge of the ethidium cation help hold the com-

pound in place. These forces of attraction explain why this compound has such a

strong association with DNA. Another example of the importance of ionic bond-

ing accounts for the source of calcium in milk. Casein, one of the major proteins

found in milk, contains phosphate anions that form ionic bonds with calcium

cations.

Sizes of the Ions

Water from many wells typically contains relatively high concentrations of

calcium, often along with some magnesium and iron. When the water is heated,

dissolved calcium bicarbonate, Ca(HCO

)

, decomposes to form rock-like de-

posits of calcium carbonate, CaCO

, also known as boiler scale. The boiler scale

Ca(HCO

)

(aq) n CaCO

(s) + H

O(l) + CO

(g)

coats the inside of the pipes that carry the water in the hot-water heater and

throughout the house, so such water is called “hard.” Hard water is not hazardous



Ca Cl



Ca Cl



2



Na Cl Na









CHC





FIGURE 8.5

Ethidium bromide. The stabilizing force

that holds the molecule in place in DNA

is the attraction between the anionic

charges on the DNA and the positive

charge on ethidium bromide.

Application

to our health. It is just not the best thing that could happen to our house-

hold pipes. If the buildup is bad enough, the pipes can become clogged,

much as heart arteries become clogged with fatty plaques, reducing blood

ﬂow in a human. Many homeowners and some cities “soften” water by

passing it through a bed of

zeolites, which exchange the calcium ions for

sodium ions that (remember our solubility rules) do not form insoluble

salts. How do these ion-exchanging zeolites work?

Zeolites like that shown in Figure 8.6 are composed of aluminum and

silicon oxide subunits containing Group IA and IIA cations. The resulting

honeycomb arrangement of subunits results in a structure full of atom-

sized holes. Ions and molecules small enough to ﬁt through these holes can

enter the zeolite, and if they are just about the same size as the holes in the

zeolite, they get stuck inside. Interactions of the ions or molecules with the

zeolite help to hold them inside. Substances that are too small easily ﬂow in

and out of the zeolite, whereas ions and molecules that are too big can’t

enter the zeolite in the ﬁrst place. Consequently, a zeolite retains only spe-

ciﬁc sizes of ions and molecules. In exchange for the trapped ions, the zeo-

lite releases ions of the same total charge. For example, when a calcium ion is

taken up by the zeolite, it typically releases two sodium ions.

By carefully constructing the zeolite, researchers have been able to develop an

“ionic sponge” that grabs only the ions of interest to them. This is vital, in the

chemical industry, to the formation of better reaction catalysts (compounds that

signiﬁcantly increase the rate of a chemical reaction without being consumed)

the ﬁltration of polluted air, and the cleanup of hazardous wastes, among other

applications. Experiments with the zeolite known as clinoptilolite, with a pore

size of 0.4 nm, indicated that it was capable of removing radioactive cesium

(

134

Cs and

137

Cs) from cows affected by the 1986 Chernobyl (Ukraine) nuclear

accident.

By far the most common use of zeolites is as an ingredient called a “builder”

in laundry detergents for the removal of calcium ions from hard water. Size is a

critical factor in the behavior of ionic compounds. Not only does size suggest

what type of zeolite can trap a particular ion, it also plays a major role in deter-

mining the structure of the ionic crystal and the strength of the ionic bond.

What contributes to the size of an ion? We can examine the electrons in an

atom more closely to determine the answer. In our previous discussion of the

shape of the orbitals (Chapter 6), we learned that the electrons penetrate deep

into the atom. Each electron helps to balance the charge of the protons in the nu-

cleus and shields the other electrons from the nuclear charge. When an electron

310 Chapter 8 Bonding Basics

“Hard” water contains relatively high concentra-

tions of calcium, typically along with magnesium

and iron. The rock-like deposits of calcium car-

bonate that form on the inside of hot water

pipes give hard water its name.

Application

HEMICAL

ENCOUNTERS:

Focus on Zeolites

Natrolite Sodalite

FIGURE 8.6

Ions can enter, and they can leave: The

structure of natrolite and sodalite. These

zeolites have pores that ions of the right

size can enter.

Visualization: Ionic Radii

is removed from an atom, the amount of shielding is reduced. We say that the ef-

fective nuclear charge experienced by each electron increases. The total number

of electron−electron repulsions decreases, so the electrons are pulled closer to

the nucleus. The net effect is a reduction in the size of the electron cloud around

the cation. The size of the cation is less than the size of the atom (see Table 8.2).

The opposite happens when an electron is added to an atom. The extra electron

increasesthe number of electron–electronrepulsions,therebyreducingtheeffective

nuclear charge felt by each electron. This causes the electron cloud to swell in size.

Theanion,then,islargerthanthe atomfromwhich itwasmade.Theseeffects continue

as more electrons are removed from or added to an atom. All of the ions shown in

Figure 8.7 have the same electron conﬁguration, but they differ in the charge of the

nucleus.We can see that as the nuclear charge increases,the size of the ion decreases.

We noted that the removal of the ﬁrst electron greatly decreases the size of an

ion. Similar reasoning holds for the removal of subsequent electrons. In general,

8.2 Ionic Bonding 311

The Radius of Some Common lons and Atoms (in picometers)

TABLE 8.2

247

Period

Group

148

265

169

227

133

186

152

IA IIA

143

117

104

2–

184

2–

140

VIA

215

113

217

135

197

2–

221

2–

198

160

113

110

3–

212

3–

171

IIIA

143

133

114

–

181

–

133

VIIA

–

216

–

195

3–

171 pm

2–

140 pm

–

133 pm

50 p

99 pm

65 pm

FIGURE 8.7

The radii of a series of isoelectronic ions.

the cation gets smaller as more electrons are removed. Removal of the ﬁrst electron

from an atom greatly decreases the radius of the resulting ion. Removal of the

second and all subsequent electrons from the ion has the same effect, but in a less

pronounced way. This is due to the fact that electrons in the same subshell do not

shield each other very effectively from the nuclear charge.

As you descend within a group in the periodic table, the principal quantum

number of the valence electrons increases. The distance from the nucleus in-

creases with an increasing principal quantum number. The size of the ion

increases. This trend is shown in Table 8.2.

EXERCISE 8.3 Ions and Their Sizes

Argon is a noble gas found in the atmosphere in relatively large concentration

(0.934% by volume). It is used as an inert atmosphere in the chemical laboratory for

reactions that would react with normal atmospheres. List four ions that are isoelec-

tronic with argon, and arrange them in order of increasing size.

First Thoughts

The question is really asking us to complete two tasks. The ﬁrst is to ﬁnd four atoms

that can be made into ions with the same electron conﬁguration as argon. There-

fore, we should limit our search to those atoms that have an atomic number close to

that of argon. Second, the question asks us to organize the ions we’ve created by size.

We know that cations are smaller than atoms and anions are larger.

Solution

 K

 Ar  Cl

−

 S

2−

Further Insight

Knowing the size and electron conﬁguration of an ion is helpful in constructing the

bigger picture for a model of a compound. The ﬁve species we’ve arranged here have

the same number of electrons around their nuclei, but they differ greatly in size.

Moreover, their having the same number of electrons doesn’t imply that these ions

have any properties in common. They are completely different in their reactivity

and in their physical properties.

PRACTICE 8.3

Arrange four ions that are isoelectronic with Ne in order of increasing size.

See Problems 23, 24, 31, and 32.

HERE’S WHAT WE KNOW SO FAR

■

The electron conﬁguration of an ion determines the number of electrons in its

outer shell.

■

We can use the electron conﬁguration of an ion to draw the Lewis dot symbol

for the ion.

■

Anions, atoms with extra electrons, are larger than their corresponding atoms.

■

Cations, atoms missing electrons, are smaller than their corresponding atoms.

■

Ionic compounds are the combination of anions and cations to make an elec-

trically neutral compound. The force of attraction between an anion and a

cation is the ionic bond.

312 Chapter 8 Bonding Basics

Energy of the Ionic Bond

According to the American Dental Association, the addition of ﬂuoride to tooth-

paste and public water systems has brought about a nationwide reduction in

tooth decay. The Centers for Disease Control suggest a “safe, effective and inex-

pensive”municipal waterway ﬂuoride concentration of between 0.7 and 1.2 parts

per million. The mineral portion of teeth is hydroxyapatite, Ca

(PO

)

(OH).

When you drink ﬂuoridated water or brush your teeth with toothpaste containing

fluoride, some of the hydroxide anions in your teeth are replaced with ﬂuoride.

(PO

)

(OH)

−−−→

(PO

)

(OH,F)

Hydroxyapatite Fluorapatite

The new mineral, called ﬂuorapatite, is much stronger than hydroxyapatite.What

accounts for the added strength? Does the size of the ﬂuoride and hydroxide ions

have something to do with the strength of the mineral?

The strength of the ionic bond is usually referred to as the

lattice enthalpy of

the ionic solid. The lattice enthalpy of a molecule is the amount of energy re-

quired to separate 1 mol of a solid ionic crystalline compound into its gaseous

ions (see the following equation). The lattice enthalpies of some common com-

pounds can be found in Table 8.3.

MX(s) n M(g)

+ X(g)

−

Although qualitative statements can be made about the relative size of the lattice

enthalpy on the basis of ionic radii, most chemists use lattice enthalpy to make

some quantitative statements about the strength of an ionic compound. Unfortu-

nately, it is quite difﬁcult to measure lattice enthalpy accurately in an ionic crys-

talline solid. However, we can calculate the lattice enthalpy using Hess’s law (see

Chapter 5) in a process known as the

Born–Haber cycle, named after two Nobel

Prize–winning German scientists (Max Born, 1882–1970, and Fritz Haber,

1868–1934).

The Born–Haber cycle is a diagrammatic representation of the formation of

an ionic crystalline solid. Figure 8.8 illustrates the Born–Haber cycle for the

formation of sodium chloride from its elements in their standard states. Al-

though chemistry is a process in which all kinds of things happen concurrently

and continuously, for clarity we often break down the Born–Haber cycle into

a series of steps.

8.2 Ionic Bonding 313

Application

HEMICAL

NCOUNTERS:

Fluoridated Water

and Tooth Decay

Lattice Enthalpies for Some Common Ionic Solids

Values in the table are in kilojoules per mole for the simple ionic compounds

(for example, the lattice energy of K

O is 2238 kJ/mol).

−

2−

1030 834 788 757 1039 2799

923 787 747 704 887 2481

821 701 682 649 789 2238

785 689 660 630 766 2163

740 659 631 604 721 ––

2913 2326 2097 1944 2870 3795

2609 2223 2132 1905 2506 3414

2341 2033 1950 1831 2141 3029

5096 4874 4711 4640 5063 13,557

5924 5376 5247 5070 5627 15,916

TABLE 8.3

314 Chapter 8 Bonding Basics

■

At the starting point to the cycle, we sublime solid sodium metal into gaseous

sodium metal. The enthalpy change for this process is known and is recorded

on the cycle.

■

Next, we dissociate the chlorine molecule into individual chlorine atoms,

using the equation that relates to the bond dissociation energy for chlorine.

Because we need only one chlorine atom, we take only half of the enthalpy

change for this process.

■

In the next steps, the gaseous atoms are ionized to the gaseous ions. Sodium

atoms are ionized to sodium cations with an enthalpy change that corresponds

to the ﬁrst ionization enthalpy. Chlorine atoms are converted into chloride

anions with an enthalpy change that corresponds to the electron afﬁnity for

chlorine.

■

In the ﬁnal step, the sodium and chloride ions are brought together to make

the ionic crystalline solid. This enthalpy change corresponds to the negative of

the lattice enthalpy for the ionic solid.

When the heat of formation for the ionic solid is known by direct measure-

ment, the direct route from the starting point to the end of the Born–Haber cycle

(the lattice enthalpy) can be calculated by summing all of the individual enthalpy

changes. Because the individual enthalpy changes are often known (heat of for-

mation, sublimation, ﬁrst ionization, bond dissociation, and electron afﬁnity),

this provides a convenient method for determining the lattice enthalpy of the

ionic crystalline solid.

Calculation of the

lattice energy (not lattice enthalpy) of an ionic compound

can also be accomplished using

Coulomb’s law. Coulomb’s law states that the force

between two particles is proportional to the product of the charges (Q) on each

particle divided by the square of their distance of separation (d). The French

physicist Charles Augustin Coulomb (1736–1806) proved this relationship by ex-

periment. Because the force of attraction between two ions is related to the lattice

energy, a modiﬁcation of Coulomb’s potential gives us the lattice energy:

Lattice energy = k



× Q

−



In this equation, k is a proportionality constant, Q

and Q

−

are the charges on the

ions, and d is the distance between the ions. The lattice energy is related to the lat-

tice enthalpy. The difference is that the lattice energy (E

lattice

) is the amount of en-

ergy released as the solid ionic crystal is formed from its gaseous ions, whereas the



sub

H = +108 kJ

IE = +496 kJ

NaCl(s)

Na(s)

Na(g)

(g)



diss

H = +122 kJ

EA = –349 kJ

U = –788 kJ



H˚



H˚ = 

sub

H + IE + 

diss

H + EA + U



H˚ = 108 + 496 + 122 – 349 – 788 = –411 kJ/mol

Cl(g)

–

(g)

FIGURE 8.8

The Born–Haber cycle for the

formation of sodium chloride.

Visualization: Born–Haber Cycle

for NaCl(

)

lattice enthalpy (H

lattice

) is the amount of heat at constant pressure necessary to

separate the solid ionic crystal into its gaseous ions. The good news is that these

two values are just about the same, though opposite in sign; lattice energy de-

scribes a release in energy (E

lattice

= “−”), whereas lattice enthalpy describes

energy that is absorbed (H

lattice

=“+”).

H

lattice

∼

−E

lattice

The magnitude of the lattice energies or lattice enthalpies increases as the

charges on the ions increase, as shown in Table 8.3. In the series ScCl

,CaCl

, and

KCl, the charges decrease on the cation: (+3), (+2), and (+1), respectively,

with a decrease in lattice enthalpies in the order ScCl

(4874 kJ/mol), CaCl

(2223 kJ/mol), KCl (701 kJ/mol). Coulomb’s law also indicates that as the

distance between the ions decreases, the lattice enthalpies increase. Evidence of

this is found by comparing LiBr (788 kJ/mol), LiCl (834 kJ/mol), and LiF

(1030 kJ/mol). The charges on the ions are the same in this series of compounds,

but the size of the anion decreases from bromide to ﬂuoride. Because the

bromide ion (Br

−

) is larger than Cl

−

or F

−

, LiBr has the smallest lattice enthalpy.

In general, lattice enthalpies are greatest for ionic compounds that are made up of

small, highly charged particles.

Let’s revisit our section-opening question of why ﬂuoride in our toothpaste is

important. As we noted then, the replacement of a hydroxy group in the hydrox-

yapatite, Ca

(PO

)

(OH), mineral that makes up our teeth gives rise to a new

mineral called ﬂuorapatite, Ca

(PO

)

(OH,F). Because F

−

is smaller than OH

−

Coulomb’s law dictates that the force of attraction between the Ca

and the F

−

should be greater than that between Ca

and OH

−

in this mineral. This is the

case with, for example, nearly all binary ionic salts of ﬂuoride compared to the

metal hydroxide, so that the lattice enthalpy of, for example, silver ﬂuoride (AgF),

which is 953 kJ/mol, is greater than the lattice enthalpy of silver hydroxide

(AgOH), 918 kJ/mol. It is therefore reasonable to consider that the lattice en-

thalpy of ﬂuorapatite is greater than that of hydroxyapatite. This is one of several

reasons why ﬂuorapatite is more stable than hydroxyapatite when bathed in our

saliva—and more resistant to the formation of cavities. Fluorapatite formation is

even more compelling when ﬂuoride is present in our mouths from municipal

water ﬂuoridation or dental treatments; this is related to a concept called chemi-

cal equilibrium, which we will consider in Chapters 16–18.

EXERCISE 8.4 Predicting Lattice Enthalpies

Use the relationship of ionic sizes to predict whether calcium ﬂuoride (found in

toothpaste) or calcium chloride (sidewalk salt) has the greater lattice enthalpy. Also

predict whether aluminum chloride or sodium chloride has the greater lattice

enthalpy.

First Thoughts

Comparing the lattice enthalpies of two ionic compounds can be accomplished by

realizing that the magnitude of the lattice enthalpy is inversely proportional to the

distance between the individual ions and directly proportional to the size of the nu-

clear charge on the ions. The distance between the ions is directly related to the

individual ionic radii.

Solution

Calcium ﬂuoride (CaF

) and calcium chloride (CaCl

) differ in the size of the anion

bound to the calcium cation. From the discussion of atomic size, we noted that

ﬂuorine is smaller than chlorine. Moreover, the radius for both anions also follows

this trend; ﬂuoride is a smaller anion than chloride. Coulomb’s law says that CaF

8.2 Ionic Bonding 315

has a larger, more positive lattice enthalpy. And indeed it does (2609 kJ/mol versus

2223 kJ/mol).

Aluminum chloride (AlCl

) and sodium chloride (NaCl) differ in the charge

of the cation. The larger charge on the aluminum cation (+3) indicates that alu-

minum chloride should have the larger lattice enthalpy. It does (5376 kJ/mol versus

787 kJ/mol).

Further Insights

Which has a greater inﬂuence on the lattice enthalpy, ionic radius or ionic charge?

The examples we’ve examined suggest that the ionic charge has a much greater

effect. We can reason, and rightly so, that the electrostatic force of attraction be-

tween a cation and an anion is a very powerful force. This powerful force assists

some proteins as they fold into a biologically active molecule.

PRACTICE 8.4

Which has the greater lattice enthalpy, FeCl

or FeCl

See Problems 25, 26, 33, 34, 100, 101, and 103.

8.3 Covalent Bonding

In addition to NaCl (table salt), sucrose (C

, common sugar) is another

compound we see at the dinner table (Figure 8.9). What is the difference between

table salt and sugar? Placed side by side, they appear relatively similar to the

naked eye. Both are colorless crystalline compounds. When added to water, they

both dissolve. However, as we saw in Chapter 4, solutions of these two com-

pounds act differently. Table salt, a strong electrolyte, dissociates into sodium

cations and chloride anions when dissolved in water. Sucrose, a nonelectrolyte,

doesn’t dissociate when it dissolves. This is why electric current passed through

the salt solution and lit the bulb, while the bulb over the sugar solution remained

unlit. Sucrose is an example of a compound containing only covalent bonds. The

atoms that make up sugar are ﬁrmly held together; bonds between these atoms

don’t dissociate upon addition to water. Because of this, an aqueous solution of

sucrose doesn’t conduct electricity. These bonds differ from the bonds in sodium

chloride, an ionic compound that dissociates in water, because the electrons in

sucrose are shared.

Knowing the type of the bonds that make up a molecule is fundamental to

those who do chemistry for a living. For example, pharmacognocists understand

that whereas many ionic compounds are readily soluble in water and can be

316 Chapter 8 Bonding Basics

FIGURE 8.9

Common table sugar. Sugar can be

crystallized in large chunks that look

much like the crystals of table salt.

Cough syrups often contain ethanol

(listed on the ingredient label as alcohol)

to increase the solubility of the active

compounds in the medicine that would

otherwise not be sufﬁciently soluble in

water.

administered to patients in aqueous solutions, many covalent compounds have

limited solubility in water. Drugs that are insoluble in water must be adminis-

tered in other solvents, which explains why cough syrups often contain ethanol.

In this section, we will examine the covalent bond up close, draw pictures of mol-

ecules that utilize this bonding scheme, and calculate the forces that hold these

bonds in place.

Description of Covalent Bonding

Opposite to the ionic end of the bonding spectrum is the covalent bond. In the

ionic bond, valence electrons gather on one of the atoms in the bond, leaving the

other atom with a deﬁciency of electrons. In the covalent bond, the atoms share

valence electrons to differing degrees. This occurs because the electronegativity of

the atoms on either end of the covalent bond are similar in magnitude, so the

atoms do not form ions. However, in order to participate in a bond and at the

same time obtain a noble gas electron conﬁguration, the atoms in the covalent

bond must share their electrons. Examples of compounds that contain covalent

bonds include H

,CH

,CO

, and H

O (Figure 8.10).

Because of the similarities in electronegativity, the majority of covalent bonds

exist between nonmetals. There are exceptions, however, such as the covalent

bonds between lead and each of the four of the carbon atoms in tetraethyl lead,

Pb(C

)

, a compound that used to be added to gasoline to improve engine

performance. Why is this an exception? The four lead–carbon covalent bonds in

the compound exist as a consequence of the relatively small difference in the

electronegativity between the lead and carbon atoms. Because the emissions-

reducing catalytic converters required on all cars and light trucks in the

United States since 1981 are destroyed by lead, and because the combustion of the

lead-containing gasoline releases the metal into the environment, this type of

automobile fuel was banned in the United States in 1986. All of the automotive

gasoline sold today in the United States is unleaded gasoline. And as a result, the

amount of lead pollution of the environment has decreased. Table 8.4 lists the

47 countries that have completely eliminated the use of leaded gasoline in motor

vehicles. A World Bank Regional Conference on “The Phase-out of Leaded Gaso-

line in Sub-Saharan Africa,” held in Senegal in 2001, set as a goal the elimination

of leaded gasoline from that entire continent by 2005. Recent United Nations and

World Bank data show that this goal has largely been achieved.

8.3 Covalent Bonding 317

FIGURE 8.10

Examples of compounds

that contain covalent

bonds include CO

, and CH

Application

HEMICAL ENCOUNTERS:

Tetraethyl Lead in

Gasoline