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

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Countries That Have Eliminated

Tetraethyl Lead Use in Motor Vehicles

Albania

Antigua

Argentina

Austria

Bahrain

Bangladesh

Belgium

Belize

Bolivia

Brazil

Canada

Colombia

Costa Rica

Denmark

Dominican

Republic

TABLE 8.4

Ecuador

Egypt

El Salvador

Finland

Germany

Guatemala

Haiti

Honduras

Hungary

Iceland

India

Ireland

Jamaica

Japan

Luxembourg

Mexico

Netherlands

New Zealand

Nicaragua

Norway

Philippines

Saudi Arabia

Singapore

Slovakia

South Korea

Sweden

Switzerland

Taiwan

Thailand

United Kingdom

United States

Vietnam

The simplest example of a compound containing a covalent bond is hydrogen

gas (H

), used in the food industry to make partially saturated oils, which contain

mostly carbon–carbon and carbon–hydrogen single bonds, as well as in the man-

ufacture of ammonia intended for agricultural use as a fertilizer. The bond be-

tween the hydrogen atoms in H

results from the sharing of one electron from

each of the atoms. The atoms on either end of the bond have the same afﬁnity for

electrons, so they must be shared if each hydrogen is to have the electron conﬁg-

uration of a noble gas. In other words, sharing the electrons between the two

atoms effectively gives each hydrogen a 1s

electron conﬁguration. The sharing of

bonding electrons is the deﬁning characteristic of a covalent bond. What holds

the two atoms together in a covalent bond? Each nucleus on either end of the

bond exerts a force of attraction on the pair of bonding electrons. This attractive

force pulls the nuclei close to form a bond. If we examine the electron cloud

around the two nuclei in molecular hydrogen, we note an interesting feature of

the covalent bond, which is shown in Figure 8.11. The density of electrons is con-

centrated between the two atoms, but a signiﬁcant amount of electron density

surrounds each of the two nuclei. As shown in Figure 8.12, the picture of electron

density in HF, a compound used to etch glass, is remarkably different because the

318 Chapter 8 Bonding Basics

H(g)H(g)H

(g)

FIGURE 8.11

Electrons hold the nuclei together in this model of hydrogen

gas. Note that the electron density, shown in red, encircles

both nuclei.

FIGURE 8.12

Compare the electron density

distribution in HF with that of H

Tetraethyl lead, Pb(C

)

, delivered from gas pumps similar to

today’s, has been banned as an additive to vehicular fuel in many

countries because of the harmful effects of lead on people and

the environment.

bonding electrons are not shared quite equally, although they are shared to an

important extent, and the bond is still considered to be covalent.

What keeps the nuclei from bumping into each other? While the attractive

force of the nuclei for the shared electrons pulls the atoms together, a repulsive

force of the like-charged nuclei pushes them apart. This repulsion keeps the nu-

clei from getting too close. It is the combination of attractive and repulsive forces

that holds the nuclei apart at a particular average distance that we refer to as the

bond length. A plot of the potential energy of two hydrogen atoms as a function

of their distance, shown in Figure 8.13, helps to illustrate this fact.

Electronegativity and the Covalent Bond

Why are electrons shared unequally in some covalent bonds such as HOF? Some

atoms have a greater attraction for shared electrons than do others. The electron

density in the covalent bond is pulled closer to the atom with the greater attrac-

tion for the bonding electrons. We say that some covalent bonds are

polarized

toward one of the atoms. In hydrogen ﬂuoride, the attraction of the ﬂuorine

atom for electrons causes the bonding electrons to spend more time at the ﬂuo-

rine end of the molecule, resulting in a net

polarization of the bond. The lack of

electron density at the hydrogen end of the bond means that some of the total nu-

clear charge isn’t balanced on the atoms, so each of the atoms in the polar cova-

lent bond possesses a partial charge. We usually indicate the charge separation

with a lowercase Greek letter delta (δ) and a plus or minus sign. We’ll discuss how

to arrive at the bonding model for HF later, but the charge separation can be

represented like this:

Methods have been developed to attempt to identify the type of bond in a

molecule on the basis of the electronegativity of the bonded atoms.

Electronega-

tivity

was deﬁned in Chapter 7 as the ability of an atom in a molecule to attract

shared electrons to itself (Figure 8.14). The atom with the stronger attraction for

the electrons pulls the bonding electrons closer to itself. By examining the differ-

ence in the electronegativity values of the two atoms involved in a bond, we can

H F

␦␦

8.3 Covalent Bonding 319

Energy (kJ/mol)

Internuclear distance (pm)

74 pm

–458

(Covalent bond

length for H—H)

FIGURE 8.13

The energy proﬁle of a covalent bond in

as a function of distance. Note how

the attractive forces and repulsive forces

balance each other at a distance of

74 pm. At this distance the interaction

has the lowest energy. To break the

covalent bond in H

and separate the

atoms would require the addition of

458 kJ/mol.

Visualization: Bonding in H

assess the polarization of the bonding electrons (also known as the bonding elec-

tron pair). For example, ﬂuorine is much more electronegative than hydrogen, so

the electrons spend more time on the ﬂuorine end of the bond in HF.

The Pauling electronegativity scale is just one of the many attempts scientists

have made to develop a general trend in the polarization of electrons in bonds.

Pauling’s electronegativity scale, the most popular among several such scales, is

based on bond energies of diatomic molecules. Other scientists have developed

useful electronegativity trends based on ionization energies and electron afﬁni-

ties (the Mulliken scale), atomic energies and covalent radii (the Allred–Rochow

scale), the “compactness of an atom’s electron cloud” (the Sanderson scale), di-

electric properties (the Phillips scale), and quantum defects (the St. John–Bloch

scale). All of these are based on the experimentally determined properties of

some standard compounds, and this makes the scales subject to some variation.

Even though all of these scales have different values for the electronegativity of

the elements, the general trends are identical among the different systems. Pauling’s

electronegativity scale, however, is the most important historically, it has proved

quite simple to use, and the conclusions drawn from it are reasonable. We will be

on safe ground using the Pauling scale.

Types of Covalent Bonding

How can we use electronegativity to determine the degree of polar character in a

bond?

In molecular hydrogen (H

), the difference in the Pauling values for elec-

tronegativity () between the two atoms is zero ( = 2.1 − 2.1 = 0). In a mole-

cule of HCl, the electronegativity difference is less than 1 electronegativity unit

( = 0.9). Between hydrogen and ﬂuorine in HF, the difference is fairly large

( = 1.9). And in table salt (NaCl), the difference in electronegativity is 2.1.

Molecular hydrogen is an example of a molecule with a covalent bond that is

not polarized. We say that it contains a

nonpolar covalent bond (  0.5).

Hydrogen chloride (HCl) and hydrogen ﬂuoride (HF), on the other hand, con-

tain fairly polarized covalent bonds. We say that these are

polar covalent bonds

(0.5    2.0). If the difference in electronegativity is very large (2.0), the

320 Chapter 8 Bonding Basics

1.1

1.4

1.3

1.2

2.1

1.0

1.5

0.9

1.2

0.8

1.0

0.8

1.0

0.7

0.9

0.7

0.9

1.3

1.2

1.1

1.5

1.4

1.3

1.1

1.3

1.6

1.5

1.6

1.8

1.7

1.5

1.9

1.8

2.2

1.9

2.2

1.9

2.2

1.9

2.4

1.6

1.7

1.9

1.6

1.7

1.8

1.5

2.0

1.8

1.9

1.8

2.5

2.0

1.9

2.1

3.0

2.4

2.1

2.0

2.5

3.5

2.8

2.5

2.2

3.0

4.0

Increasing electronegativity

Decreasing electronegativity

Rf Db Sb Bh Hs Mt Uun Uuu Uub

FIGURE 8.14

Pauling’s electronegativity values for the

elements of the periodic table.

Compare the electron density

distributions of H

, HF, and HCl

with their electronegativity

values.

more electronegative atom’s ability to attract electrons in the bond overcomes the

tendency to share electrons, and we call this an ionic bond. The electronegativity

difference between sodium and chlorine in table salt ( =2.1) is characteristic of

an ionic bond. The difference in electronegativity can be useful in determining

the type of bond in a compound, but the values shown here are useful guidelines

only for the determination of the bonding pattern. We can say this in another

way: There is no ﬁxed cutoff. For example, it is not chemically reasonable to say

that two bonds with electronegativity differences that are within, say, 0.1 of each

other are signiﬁcantly different in bonding character. Electronegativity differences

describe a continuum, not a discrete “on-or-off,” as with a light switch.

Let’s revisit our example of a covalent bond between a metal and a nonmetal

in tetraethyl lead. Does the electronegativity difference between the lead and car-

bon atoms indicate that the bond should be covalent? Examination of the table of

electronegativity values indicates that the difference is only 0.6 electronegativity

unit. On the basis of this information, we can say that the PbOC bonds in

tetraethyl lead are polar covalent bonds and that these bonds are not as polarized

as the bond in HF or HCl. We’ll discuss the implications of this fact in the next

section.

EXERCISE 8.5 What Type of Bond Is It?

For each of these compounds, indicate whether the bond is an ionic bond, a polar

covalent bond, or a nonpolar covalent bond.

a. KCl (a replacement for table salt in low-sodium diets)

b. H

O (water)

c. Br

(used in the manufacture of brominated vegetable oils)

First Thoughts

To answer this question, we need to consider the electronegativity difference

between the bonded atoms or ions.

Solution

a. We’d predict an ionic bond between potassium and chloride. The two atoms

differ by 2.2 electronegativity units.

b. A polar covalent bonding pattern is predicted (the difference is 1.4

electronegativity units). Hydrogen is less electronegative than oxygen, so the

majority of the electron density in the bonds lies closer to the oxygen.

c. A nonpolar covalent bond is indicated here (no difference in electronegativity).

Further Insights

Differences in electronegativity can be used to identify the type of bond in a com-

pound. However, is it correct to consider that NaCl ( = 2.1) contains a purely

ionic bond and that NaBr ( =1.9) contains a polar covalent bond? What do we say

about the bond in LiCl ( =2.0)? Questions such as these help illustrate that bond-

ing is a spectrum, from the “purely”ionic to the “purely” nonpolar covalent. Almost

all bonds have some ionic and some covalent character. We’d surmise that the bond

in NaCl is more ionic than the bond in either LiCl or NaBr, but we’d realize that it,

too, has some covalent character.

PRACTICE 8.5

Predict the type of bond (covalent, polar covalent, or ionic) present in each of these

binary compounds: NO, F

,MgO.

See Problems 51 and 52.

8.3 Covalent Bonding 321

Here is the electron density distribution

in tetraethyl lead. How does it compare

to those of other covalent molecules?

Modeling a Covalent Bond—Lewis Structures

Felix Hoffman, a chemist hired by the Friedrich Bayer & Co. fabric dye plant, was

charged in the 1890s with ﬁnding some better products that could be made by the

company. One of the ideas that came to Hoffman was related to his father’s

rheumatism. At the time, this ailment and other pains were treated with large

doses of salicylic acid, found in the inner bark of several types of willow trees.

Unfortunately, Hoffman’s father was unable to take this pain medication because

its acidity irritated his stomach and throat. In 1897, Hoffman set about trying

to reduce the acidity of the medicine. To do this, he needed to know how the

arrangement of the atoms made the compound acidic. After some experimenta-

tion, Hoffman was able to produce a compound that reduced this acidity. His

product, acetylsalicylic acid (C

) later named aspirin, reduced the harmful

irritation by chemically modifying one of the acidic parts of salicylic acid to make

it less acidic. Knowing how the atoms are attached, and what kind of bonds link

the atoms, is important to understanding how compounds react and interact in

the body. One of the ways to illustrate how atoms are attached in covalent bonds

is to build a model of the molecule using Lewis dot structures. After we discuss

the basic ideas, we will draw the Lewis dot structure of aspirin.

Rules describing how the atoms and electrons are placed in a Lewis dot struc-

ture enable us to draw compounds in a systematic way. These rules, found in

Table 8.5, include drawing a skeletal picture of the molecule and then placing

extra valence electrons in the skeleton until the model of the compound is com-

plete. We will use these rules as we draw a Lewis dot structure of molecular

hydrogen (H

According to Table 8.5, the ﬁrst step is to count all of the valence electrons on

the atoms in the molecule. Because valence electrons are involved in bonding,

knowing the total number of these electrons helps us determine the number of

bonding pairs (electrons involved in bonding) and the number of lone pairs (elec-

trons not involved in bonding). In hydrogen, there are two total valence electrons

(one from each atom). In step 2 (Table 8.5), we draw a skeleton for the molecule.

The guidelines (steps 2a–e) for drawing a skeletal picture of a molecule are based

on preferences that atoms have for particular locations in a molecule. There are

only two atoms in the molecule, so the skeleton of H

322 Chapter 8 Bonding Basics

Application

A Set of Rules for Drawing Lewis Dot Structures

1. Determine the total number of valence electrons.

2. Determine the skeletal structure.

a. Hydrogen atoms are on the edges of the molecule.

b. The central atom has the lowest electronegativity. (There are many exceptions.)

c. In oxoacids, hydrogens are usually on the oxygens.

d. Think compact and symmetric.

e. Use intuition.

3. Draw the Lewis dot structure.

a. Draw the bonds connecting the atoms.

b. Determine the number of electrons remaining.

c. Place the remaining electrons as lone pairs, beginning on the most

electronegative atoms, until each atom has an octet.

d. All remaining lone pairs go on the central atom.

e. Assign formal charges, and redraw bonding electrons if necessary.

TABLE 8.5

Video Lesson: Lewis Dot

Structures for Covalent Bonds

Video Lesson: Predicting Lewis

Dot Structures

In step 3a, we place pairs of electrons to indicate bonding pairs between each

atom. For molecular hydrogen, the electrons are placed as follows;

H • • H

Most often, we use shorthand to show bonding pairs of electrons by replacing the

bonding electrons with a line. We draw the line to show that the atoms are chem-

ically bonded together in the molecule.

HOH

In steps 3b and 3c, we place all remaining electrons around the more electroneg-

ative atom ﬁrst until the octet rule (or duet rule, for hydrogen) is satisﬁed.

Because we do not have any remaining electrons, and because the duet rule is

satisﬁed for both atoms, we are ﬁnished.

Hydrogen ﬂuoride is either a fuming gas or a liquid, depending on the

temperature of use (its boiling point is 19.5

C, about 67

F), and has a host of

industrial applications. For example, it is used as a raw material in the production

of chloroﬂuorohydrocarbons (CFCs), insecticides, and fertilizers; as a catalyst in

the production of pharmaceuticals; in the manufacture of semiconductors; and

in etching glass (see Figure 8.15). Let’s use the rules in Table 8.5 to prepare a Lewis

dot structure for HF. Step 1 requires us to count the valence electrons on all the

atoms. Hydrogen has one valence electron, and ﬂuorine has seven, so we have a

total of eight valence electrons with which to work. Next, in step 2, we draw the

best skeletal structure of the molecule:

There are only two atoms in the molecule, so we’ll draw them next to each other.

Our bonding pair of electrons is represented as a line connecting the two atoms

(step 3a).

HOF

We’ve used two electrons in our skeleton to represent the bond between the two

atoms, so we have six electrons remaining (step 3b). These are placed in pairs

around the more electronegative atom until the octet rule (or duet rule) is satis-

ﬁed (step 3c). If any electrons remain (step 3d), the pairs are placed around the

other atoms.

Because we used all of the electrons to fulﬁll the octet rule around ﬂuorine, and

because both atoms now satisfy the octet rule (or duet rule), we have completed

the structure. Both ﬂuorine and hydrogen in HF have electron conﬁgurations of

a noble gas.

Formal Charges

Methanol (also known as methyl alcohol and wood alcohol, CH

O) is a com-

pound being advanced as a potential substitute for gasoline, because it has a high-

octane-rating equivalent. Methanol is also a renewable resource obtained from

the fermentation of cellulose-containing materials (such as wood). The following

three structures satisfy the octet rule (or duet rule) for every atom. Only one

of them is methanol.

H H

H F

8.3 Covalent Bonding 323

Application

FIGURE 8.15

This punch bowl (American, 1918–1919)

was made from blown glass. The pattern

was applied by etching the glass using a

dilute hydrogen ﬂuoride solution.

Which structure is it? Chemists use formal charges on atoms as one important

piece of evidence in determining the most reasonable structure for a compound

in which more than one structure might be possible.

What is a formal charge (see step 3e in Table 8.5) and how does it help us to

select the most reasonable structure? The

formal charge on an atom is the differ-

ence between the number of valence electrons on the free atom and the number

of electrons assigned to the atom when it is part of a molecule. The use of formal

charges on atoms indicates that the atoms exhibit an imbalance between the

number of electrons and the number of protons on the atom. Although this im-

balance can be estimated by assigning an oxidation number to each of the atoms

(see Chapter 4), the method of calculating formal charges lends us additional

guidance. Oxidation numbers provide an indication of the charge an atom would

have if it were completely ionic. Formal charges, on the other hand, provide the

charge on an atom assuming there is no difference in electronegativity among

the atoms in a structure. In other words, formal charges assume that electrons

are shared equally, as is the case in the covalent bonding pattern.

Formal charge = valence electrons − # bonds − # nonbonded electrons

Mathematically, the formal charge equals the number of valence electrons on

the free atom minus the number of bonds to that atom minus the number of

nonbonded electrons. As a check of our math, the sum of the formal charges on

all the atoms should equal the total charge of the molecule (0) or ion (+ or −).

We can use this information to calculate the formal charge on each atom in a sim-

ple molecule, our Lewis dot structure of HF. The ﬂuorine atom has seven valence

electrons (in Group VIIA, noted from the periodic table). Subtracting the num-

ber of bonds in the structure (one) and the number of nonbonded electrons (six)

from this number gives a formal charge of zero for the atom. Similar calculations

can be done with the hydrogen atom, which also has a formal charge of zero.

When we draw Lewis dot structures, the best structure is one that satisﬁes the

largest number of formal charge rules. The best structure:

■

has the smallest magnitude for all of the formal charges

■

places negative formal charges on the more electronegative atoms

■

has the smallest number of nonzero formal charges

We can now consider the second part of our section-opening question:

Why is

it useful to know the formal charges on atoms within a particular Lewis dot struc-

ture?

Let’s return to our discussion of methanol. Calculating the formal charges

on the central atoms in each of the possible structures enables us to choose the

one on the left as the correct structure. Every atom in both structures has an octet

of electrons, but only the structure on the left shows a formal charge of zero on

each atom. This is one piece of evidence that the structure on the left is likely to

be the most energetically stable structure of the three.

All zero formal charges

Methanol

Formal charge on C  1

Formal charge on O  1

Not methanol

Formal charge on C  2

Formal charge on O  2

Not methanol

324 Chapter 8 Bonding Basics

Video Lesson: Formal Charge

EXERCISE 8.6 Return to Basics

Pioneers in the American West made most of their everyday items from natural

sources. On the treeless plains of the Midwest they built homes of sod, burned dried

buffalo dung in the stove for heat, and made lye soap. Lye (a mixture of sodium hy-

droxide, NaOH, and potassium hydroxide, KOH) obtained from ﬁreplace ashes was

used to make soap from animal fat. What is the Lewis dot structure model for the

hydroxide ion (OH

−

)? On which atom in lye does the nonzero formal charge reside?

Solution

The hydroxide ion can best be modeled by drawing the skeleton of the molecule

with a single bond between the oxygen and the hydrogen. Placing the remaining

electrons in pairs around the most electronegative element gives the structure

shown below. The brackets indicate that the negative charge has not been assigned

to a particular atom.

Formal charge (F.C.) calculations indicate that the charge must reside on the oxygen

(F.C. = 6 − 1 − 6 =−1) and not on the hydrogen (F.C. = 1 −1 − 0 = 0), which

makes sense because the oxygen atom has a much higher electronegativity (3.5)

than the hydrogen atom (2.1). Our structure can be drawn as follows:

PRACTICE 8.6

What is the formal charge on each of the atoms in OCl

−

? in H

CONH

See Problems 55 and 56.

Multiple Bonds

We often discuss carbon dioxide in this text because it is important to life in pho-

tosynthesis and cell respiration, and it is one product of combustion. The Pauling

electronegativity values of the atoms indicate that CO

probably contains polar

covalent bonds. Following our guidelines for constructing a Lewis dot structure

of CO

, we count the valence electrons in all atoms (4 in carbon and 6 in each of

the two oxygen atoms = 16 valence electrons) and build the skeleton. The less

electronegative carbon atom is the central atom in this molecule, and the oxygen

atoms are symmetrically placed about the carbon. Next, the bonding pairs of

electrons are placed between the atoms. The remaining 12 electrons are then

placed on the more electronegative oxygen atoms to ﬁll their octets. The result is

shown below. Is this a satisfactory Lewis dot structure?

No. Both oxygen atoms have octets of electrons, but the carbon has only four

electrons (all bonding electrons) around it. What are the formal charges on each

atom in the structure?

1Formal charges

Formal charge (O)  6  1  6  1

Formal charge (C)  4  2  0  2

12

O O

OCO



8.3 Covalent Bonding 325

Our calculations reveal a lot of nonzero formal charges. Although the formal

charges add up to zero, the size and number of formal charges indicate that some-

thing may be amiss. Because energy is required to cause a separation of charges

(as indicated by the presence of nonzero formal charges in a molecule), the bet-

ter structure is generally that which minimizes formal charges. You also might

have noted that the octet rule is disobeyed for carbon in the structure. To elimi-

nate the charge on the oxygen atoms, let’s move a lone pair from each oxygen and

place them as a bond between the oxygen and the adjacent carbon. These are

called double bonds.

The resulting model has no nonzero formal charges and allows each atom to

satisfy the octet rule. Overall, the molecule is neutral. How is this model different

from the structures we discussed previously? Does the existence of more than one

pair of electrons between two atoms indicate something about the properties of

this molecule? We’ll discuss the answers to these questions in the next section.

The best sign that we’ve constructed a good model is its agreement with observed

properties for the compound. In the case of CO

, the Lewis dot structure model

does agree with the experimentally determined shape and polarity of carbon

dioxide. Based on differences in electronegativity, we observe that the electrons

are polarized toward each end of the molecule.

Multiple covalent bonds occur in many molecules. Carbon dioxide has two

double bonds. Each oxygen atom shares four bonding electrons with a carbon

atom. In a molecule of nitrogen, three pairs of electrons are used to satisfy the

octet rule for each of the atoms. The resulting molecule contains two nitrogen

atoms that share six electrons. A

triple bond links the atoms. Nitrogen, which

makes up nearly 80% of the air we breathe, is a very stable molecule that contains

a particularly strong triple bond, with the Lewis structure shown in Figure 8.16.

EXERCISE 8.7 Chemical Warfare and Bonding

Phosgene (COCl

), a highly toxic gas, was used in World War I and several subse-

quent wars to kill soldiers who were hiding in places that bombs couldn’t penetrate.

The gas, now used in industry to make polymers, pharmaceuticals, herbicides,

and other useful compounds, reacts with water and other electron-rich molecules

(molecules with lone pairs of electrons). Draw the best Lewis dot structure for this

molecule.

First Thoughts

All of the bonds in phosgene are covalent (COCl

contains only nonmetals), but a

Lewis dot structure model of the compound may help explain why phosgene reacts

with electron-rich molecules.

Solution

The skeletal picture of the molecule places the carbon in the center (least elec-

tronegative). The other atoms are placed symmetrically about the carbon. Of the

24 total electrons, 6 (three pairs) are used to connect the atoms. The remaining

18 electrons (nine pairs) are placed as lone pairs around the more electronegative

atoms. The result is that all of the atoms, except the carbon, have a full octet. We

need to share a lone pair of electrons with carbon to satisfy the octet rule on each

atom. Calculation of the formal charges on each of the atoms indicates that the oxy-

gen and the carbon atoms should share another pair of electrons.

0Formal charges 00

O O

326 Chapter 8 Bonding Basics

FIGURE 8.16

(Top) The double bonds in carbon

dioxide (CO

). (Bottom) The triple

bond in nitrogen (N

The best model of this molecule therefore shows a double bond between the carbon

and the oxygen atoms and single bonds between the carbon and the chlorine atoms.

We could have drawn a Lewis dot structure that placed a double bond between the

carbon and one of the chlorine atoms. This would have changed the formal charge

on carbon to zero, but the formal charge on the chlorine atom would become

+1,

and the formal charge on oxygen would remain −1. The resulting structure would

not be the best Lewis structure because of the existence of nonzero formal charges.

Further Insights

Does this explain why phosgene is so reactive? By noting how the bonds in the mol-

ecule are polarized, we can identify why this molecule reacts with electron-rich

compounds. According to the electronegativity values of the atoms, each of the

bonds is polarized away from the carbon atom. Electron-rich molecules such as

water (H

O), ammonia (NH

), and other compounds containing lone pairs can

react with the electron-starved carbon atoms in phosgene.

PRACTICE 8.7

Draw the Lewis dot structure for C

. for CH

See Problems 39–48, 61–66, and 99.

Resonance Structures

Carbonate ion (CO

2−

) is a common polyatomic ion found in limestone, baking

powder, and baking soda. Addition of acid to the carbonate ion causes the for-

mation of carbonic acid (H

), which decomposes rapidly into water (H

and carbon dioxide (CO

). In baking, the carbon dioxide that is released causes

the bread to rise and makes its texture lighter.

Our ﬁrst attempt at drawing the Lewis dot structure of the carbonate ion

results in the structure shown below. Carbonate has 24 electrons, 2 of them

responsible for the −2 charge, probably electrons from calcium (CaCO

), sodium

(Na

), or whatever salt resulted in a cation that donated electrons to the

carbonate anion. The carbon atom in our structure still needs to share electrons

to satisfy the octet rule. Which atom is most likely involved in sharing electrons?

Using the formal charges on the atoms, we could reconﬁgure our electrons to

participate in a double bond with the carbon. At this point, the positive charge on

the carbon atom is gone, and all of the valences are ﬁlled (the octet rule is satis-

ﬁed). The sum of the formal charges is equivalent to the charge on the carbonate

ion. This is a good Lewis dot structure for carbonate.

11

1

11

1

Cl ClC

F.C. (O)  6  2  4  0

F.C. (C)  4  4  0  0

F.C. (Cl)  7  1  6  0

Cl ClC

F.C. (O)  6  1  6  1

F.C. (C)  4  3  0  1

F.C. (Cl)  7  1  6  0

8.3 Covalent Bonding

327

Video Lesson: Resonance

Structures