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370 Chapter 12 Chemical Bonding



12.6 Lewis Structures

OBJECTIVE:



Courtesy of the University Archives/Bancroft Library/University of California,

Berkeley #UARC PIC 13:596



Remember that the electrons in

the highest principal energy

level of an atom are called the

valence electrons.



G. N. Lewis in his lab.



Module 12: Drawing Lewis

Electron Dot Structures covers

concepts in this section.



To learn to write Lewis structures.

Bonding involves just the valence electrons of atoms. Valence electrons are transferred when a metal and a nonmetal react to form an ionic compound.

Valence electrons are shared between nonmetals in covalent bonds.

The Lewis structure is a representation of a molecule that shows how

the valence electrons are arranged among the atoms in the molecule. These

representations are named after G. N. Lewis, who conceived the idea while

lecturing to a class of general chemistry students in 1902. The rules for writing Lewis structures are based on observations of many molecules from which

chemists have learned that the most important requirement for the formation of

a stable compound is that the atoms achieve noble gas electron configurations.

We have already seen this rule operate in the reaction of metals and

nonmetals to form binary ionic compounds. An example is the formation of

KBr, where the Kϩ ion has the [Ar] electron configuration and the BrϪ ion has

the [Kr] electron configuration. In writing Lewis structures, we include only

the valence electrons. Using dots to represent valence electrons, we write the

Lewis structure for KBr as follows:

Kϩ



[ Br ]Ϫ



Noble gas

configuration [Ar]



Noble gas

configuration [Kr]



No dots are shown on the Kϩ ion because it has lost its only valence electron

(the 4s electron). The BrϪ ion is shown with eight electrons because it has a

filled valence shell.

Next we will consider Lewis structures for molecules with covalent

bonds, involving nonmetals in the first and second periods. The principle

of achieving a noble gas electron configuration applies to these elements as

follows:

1. Hydrogen forms stable molecules where it shares two electrons. That

is, it follows a duet rule. For example, when two hydrogen atoms,

each with one electron, combine to form the H2 molecule, we have

H



H



H H

By sharing electrons, each hydrogen in H2 has, in effect, two

electrons; that is, each hydrogen has a filled valence shell.

H

1s

H2

[He] configuration



H

1s



2. Helium does not form bonds because its valence orbital is already

filled; it is a noble gas. Helium has the electron configuration 1s2

and can be represented by the Lewis structure

He

[He] configuration



12.6 Lewis Structures



Carbon, nitrogen, oxygen, and

fluorine almost always obey the

octet rule in stable molecules.



371



3. The second-row nonmetals carbon through fluorine form stable

molecules when they are surrounded by enough electrons to fill

the valence orbitals—that is, the one 2s and the three 2p orbitals.

Eight electrons are required to fill these orbitals, so these elements

typically obey the octet rule; they are surrounded by eight

electrons. An example is the F2 molecule, which has the following

Lewis structure:

F ⎯⎯⎯⎯⎯⎯→ F F ←⎯⎯⎯⎯⎯⎯ F

F atom with seven

valence electrons



F2

molecule



F atom with seven

valence electrons



Note that each fluorine atom in F2 is, in effect, surrounded by eight

valence electrons, two of which are shared with the other atom.

This is a bonding pair of electrons, as we discussed earlier. Each

fluorine atom also has three pairs of electrons that are not involved

in bonding. These are called lone pairs or unshared pairs.

4. Neon does not form bonds because it already has an octet of

valence electrons (it is a noble gas). The Lewis structure is

Ne

Note that only the valence electrons (2s22p6) of the neon atom

are represented by the Lewis structure. The 1s2 electrons are core

electrons and are not shown.

Lewis structures show only

valence electrons.



Next we want to develop some general procedures for writing Lewis

structures for molecules. Remember that Lewis structures involve only the

valence electrons of atoms, so before we proceed, we will review the relationship of an element’s position on the periodic table to the number of valence electrons it has. Recall that the group number gives the total number

of valence electrons. For example, all Group 6 elements have six valence

electrons (valence configuration ns2np4).



Group 6

O

2s22p4

Group

6



S

3s23p4



Se

4s24p4



Te

5s25p4



372 Chapter 12 Chemical Bonding

Similarly, all Group 7 elements have seven valence electrons (valence configuration ns2np5).

Group 7

F

2s22p5

Group

7



Cl

3s23p5



Br

4s24p5



I

5s25p5



In writing the Lewis structure for a molecule, we need to keep the following

things in mind:

1. We must include all the valence electrons from all atoms. The total

number of electrons available is the sum of all the valence electrons

from all the atoms in the molecule.

2. Atoms that are bonded to each other share one or more pairs of

electrons.

3. The electrons are arranged so that each atom is surrounded by

enough electrons to fill the valence orbitals of that atom. This

means two electrons for hydrogen and eight electrons for secondrow nonmetals.

The best way to make sure we arrive at the correct Lewis structure for a

molecule is to use a systematic approach. We will use the approach summarized by the following rules.



Steps for Writing Lewis Structures

Step 1 Obtain the sum of the valence electrons from all of the atoms. Do not

worry about keeping track of which electrons come from which

atoms. It is the total number of valence electrons that is important.

Step 2 Use one pair of electrons to form a bond between each pair of bound

atoms. For convenience, a line (instead of a pair of dots) is often used

to indicate each pair of bonding electrons.

Step 3 Arrange the remaining electrons to satisfy the duet rule for hydrogen

and the octet rule for each second-row element.

To see how these rules are applied, we will write the Lewis structures of

several molecules.



C H E M I S T R Y I N F OCUS



O



ne of the problems we face in modern society

is how to detect illicit substances, such as drugs

and explosives, in a convenient, accurate manner. Trained dogs are often used for this purpose

because of their acute sense of smell. Now several researches are trying to determine whether

insects, such as honeybees and wasps, can be

even more effective chemical detectors. In fact,

studies have shown that bees can be trained in

just a few minutes to detect the smell of almost

any chemical.

Scientists at Los Alamos National Laboratory

in New Mexico are designing a portable device

using bees that possibly could be used to sniff out

drugs and bombs at airports, border crossings,

and schools. They call their study the Stealthy Insect Sensor Project. The Los Alamos project is

based on the idea that bees can be trained to associate the smell of a particular chemical with a

sugary treat. Bees stick out their “tongues” when

they detect a food source. By pairing a drop of

sugar water with the scent of TNT (trinitrotoluene) or C-4 (composition 4) plastic explosive

about six times, the bees can be trained to extend

their proboscis at a whiff of the chemical alone.

The bee bomb detector is about half the size of a

shoe box and weighs 4 lb. Inside the box, bees are

lined up in a row and strapped into straw-like

tubes, then exposed to puffs of air as a camera

monitors their reactions. The signals from the

video camera are sent to a computer, which analyzes the bees’ behavior and signals when the

bees respond to the particular scent they have

been trained to detect.

A project at the University of Georgia uses

tiny parasitic wasps as a chemical detector. Wasps



EXAMPLE 12.2



do not extend their tongues when they detect a

scent. Instead, they communicate the discovery of

a scent by body movements that the scientists call

“dances.” The device, called the Wasp Hound,

contains a team of wasps in a hand-held ventilated cartridge that has a fan at one end to draw

in air from outside. If the scent is one the wasps

do not recognize, they continue flying randomly.

However, if the scent is one the wasps have been

conditioned to recognize, they cluster around the

opening. A video camera paired with a computer

analyzes their behavior and signals when a scent

is detected.

The insect sensors are now undergoing field

trials, which typically compare the effectiveness

of insects to that of trained dogs. Initial results

appear promising, but the effectiveness of these

devices remains to be proved.



Los Alamos National Laboratory. Photo by Leroy Sanchez



To Bee or Not to Bee



A honeybee receives a fragrant reminder of its

target scent each morning and responds by

sticking out its proboscis.



Writing Lewis Structures: Simple Molecules

Write the Lewis structure of the water molecule.

SOLUTION

We will follow the steps listed on page 372.



373



374 Chapter 12 Chemical Bonding

Step 1 Find the sum of the valence electrons for H2O.

1



ϩ



c

H

(Group 1)



ϩ



1



ϭ 8 valence electrons



6



c

H

(Group 1)



c

O

(Group 6)



Step 2 Using a pair of electrons per bond, we draw in the two OOH bonds,

using a line to indicate each pair of bonding electrons.

HOOOH

Note that

HOOOH represents H O H

Step 3 We arrange the remaining electrons around the atoms to achieve a

noble gas electron configuration for each atom. Four electrons have been

used in forming the two bonds, so four electrons (8 Ϫ 4) remain to be distributed. Each hydrogen is satisfied with two electrons (duet rule), but oxygen needs eight electrons to have a noble gas electron configuration. So the

remaining four electrons are added to oxygen as two lone pairs. Dots are used

to represent the lone pairs.

H

H



might also be drawn as

H O H



H Lone pairs



O



H



O



H

→



This is the correct Lewis structure for the water molecule. Each hydrogen

shares two electrons, and the oxygen has four electrons and shares four to

give a total of eight.



→



O



→



H



2eϪ 8eϪ 2eϪ

Note that a line is used to represent a shared pair of electrons (bonding electrons) and dots are used to represent unshared pairs.



Self-Check EXERCISE 12.2 Write the Lewis structure for HCl.

See Problems 12.59 through 12.62. ■



12.7 Lewis Structures of Molecules

with Multiple Bonds

OBJECTIVE:



To learn how to write Lewis structures for molecules with multiple bonds.

Now let’s write the Lewis structure for carbon dioxide.

Step 1 Summing the valence electrons gives

4

c

C

(Group 4)



ϩ



6

c

O

(Group 6)



ϩ



6

c

O

(Group 6)



ϭ 16



C H E M I S T R Y I N F OCUS

Hiding Carbon Dioxide



As we discussed in Chapter 11 (see ”Chemistry



The injection of CO2 into the earth’s crust is

already being undertaken by various oil companies. Since 1996, the Norwegian oil company

Statoil has separated more than 1 million tons of

CO2 annually from natural gas and pumped it

into a saltwater aquifer beneath the floor of the

North Sea. In western Canada a group of oil companies has injected CO2 from a North Dakota synthetic fuels plant into oil fields in an effort to increase oil recovery. The oil companies expect to

store 22 million tons of CO2 there and to produce

130 million barrels of oil over the next 20 years.

Sequestration of CO2 has great potential as

one method for decreasing the rate of global

warming. Only time will tell whether it will work.



in Focus: Atmospheric Effects,” page 326), global

warming seems to be a reality. At the heart of this

issue is the carbon dioxide produced by society’s

widespread use of fossil fuels. For example, in the

United States, CO2 makes up 81% of greenhouse

gas emissions. Thirty percent of this CO2 comes

from coal-fired power plants used to produce

electricity. One way to solve this problem would

be to phase out coal-fired power plants. However, this outcome is not likely because the

United States possesses so much coal (at least a

250-year supply) and coal

is so cheap (about $0.03

per pound). Recognizing

this fact, the U.S. government has instituted a research program to see

if the CO2 produced at

CO2 stored in geologic disposal

power plants can be captured and sequestered

(stored) underground in

deep geological formations. The factors that

Unmineable

need to be explored to

coal beds

Enhanced

oil recovery

determine whether seDepleted oil

questration is feasible are

or gas reserves

the capacities of underground storage sites and

the chances that the sites

Deep saline formation

will leak.



O



C



O



Step 2 Form a bond between the carbon and each oxygen:

OOCOO



represents

O C O



O



CO2 capture at

power stations



C



represents

O C O



O



Step 3 Next, distribute the remaining electrons to achieve noble gas electron configurations on each atom. In this case twelve electrons (16 Ϫ 4) remain after the bonds are drawn. The distribution of these electrons is determined by a trial-and-error process. We have six pairs of electrons to

distribute. Suppose we try three pairs on each oxygen to give

O



C



O



Is this correct? To answer this question we need to check two things:



375

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376 Chapter 12 Chemical Bonding

1. The total number of electrons. There are sixteen valence electrons in

this structure, which is the correct number.

2. The octet rule for each atom. Each oxygen has eight electrons

around it, but the carbon has only four. This cannot be the correct

Lewis structure.

How can we arrange the sixteen available electrons to achieve an octet

for each atom? Suppose we place two shared pairs between the carbon and

each oxygen:

O



represents



C



O



Now each atom is surrounded by eight electrons, and the total number

of electrons is sixteen, as required. This is the correct Lewis structure for carbon dioxide, which has two double bonds. A single bond involves two

atoms sharing one electron pair. A double bond involves two atoms sharing two pairs of electrons.

In considering the Lewis structure for CO2, you may have come up with

O



represents

O CO



O



8

8

8

electrons electrons electrons



O C O



O



C



→



O



→



C



→



O



C



O



or



O



C



O



Note that both of these structures have the required sixteen electrons

and that both have octets of electrons around each atom (verify this for

yourself). Both of these structures have a triple bond in which three electron pairs are shared. Are these valid Lewis structures for CO2? Yes. So there

really are three Lewis structures for CO2:

O



C



O



O



C



O



O



C



O



This brings us to a new term, resonance. A molecule shows resonance when

more than one Lewis structure can be drawn for the molecule. In such a case we

call the various Lewis structures resonance structures.

Of the three resonance structures for CO2 shown above, the one in the

center with two double bonds most closely fits our experimental information about the CO2 molecule. In this text we will not be concerned about

how to choose which resonance structure for a molecule gives the “best” description of that molecule’s properties.

Next let’s consider the Lewis structure of the CNϪ (cyanide) ion.

Step 1 Summing the valence electrons, we have

CNϪ

4 ϩ 5 ϩ 1 ϭ 10

Note that the negative charge means an extra electron must be added.

Step 2 Draw a single bond (CON).

Step 3 Next, we distribute the remaining electrons to achieve a noble gas

configuration for each atom. Eight electrons remain to be distributed. We

can try various possibilities, such as

C



N



or



C



N



or



C



N



These structures are incorrect. To show why none is a valid Lewis structure,

count the electrons around the C and N atoms. In the left structure, neither



C H E M I S T R Y I N F OCUS



Eating the right foods is critical to our health. In

particular, certain vegetables, although they do

not enjoy a very jazzy image, seem especially important. A case in point is broccoli, a vegetable

with a humble reputation that packs a powerful

chemistry wallop.

Broccoli contains a chemical called sulforaphane, which has the following Lewis structure:

CH3



S



(CH2)4



N



C



S



O



Experiments indicate that sulforaphane furnishes

protection against certain cancers by increasing

the production of enzymes (called phase 2 enzymes) that “mop up” reactive molecules that

can harm DNA. Sulforaphane also seems to combat bacteria. For example, among the most common harmful bacteria in humans is Helicobacter

pylori (H. pylori), which has been implicated in

the development of several diseases of the stomach, including inflammation, cancer, and ulcers.

Antibiotics are clearly the best treatment for H.

pylori infections. However, especially in developing countries, where H. pylori is rampant, antibi-



C



N



otics are often too expensive to be available to

the general population. In addition, the bacteria

sometimes evade antibiotics by “hiding” in cells

on the stomach walls and then reemerging after

treatment ends.

Studies at Johns Hopkins in Baltimore and

Vandoeuvre-les Nancy in France have shown that

sulforaphane kills H. pylori (even when it has

taken refuge in stomach-wall cells) at concentrations that are achievable by eating broccoli. The

scientists at Johns Hopkins also found that sulforaphane seems to inhibit stomach cancer in

mice. Although there are no guarantees that

broccoli will keep you healthy, it might not hurt

to add it to your diet.



atom satisfies the octet rule. In the center structure, C has eight electrons but

N has only four. In the right structure, the opposite is true. Remember that

both atoms must simultaneously satisfy the octet rule. Therefore, the correct

arrangement is

C



represents

C



Squared Studio/PhotoDisc/Getty Images



Broccoli—Miracle Food?



N



N



(Satisfy yourself that both carbon and nitrogen have eight electrons.) In this

case we have a triple bond between C and N, in which three electron pairs

are shared. Because this is an anion, we indicate the charge outside of square

brackets around the Lewis structure.

[ C



N ]Ϫ



In summary, sometimes we need double or triple bonds to satisfy the

octet rule. Writing Lewis structures is a trial-and-error process. Start with

single bonds between the bonded atoms and add multiple bonds as

needed.

We will write the Lewis structure for NO2Ϫ in Example 12.3 to make sure

the procedures for writing Lewis structures are clear.



377



378 Chapter 12 Chemical Bonding

EXAMPLE 12.3



Writing Lewis Structures: Resonance Structures

Write the Lewis structure for the NO2Ϫ anion.

SOLUTION

Step 1 Sum the valence electrons for NO2Ϫ.

Valence electrons: 6 ϩ 5 ϩ 6 ϩ

O



N



O



1

Ϫ1

charge



ϭ 18 electrons



Step 2 Put in single bonds.

OONOO

Step 3 Satisfy the octet rule. In placing the electrons, we find there are two

Lewis structures that satisfy the octet rule:

[O



N



O ]Ϫ



and



[ O



N



O ]Ϫ



Verify that each atom in these structures is surrounded by an octet of electrons. Try some other arrangements to see whether other structures exist in

which the eighteen electrons can be used to satisfy the octet rule. It turns out

that these are the only two that work. Note that this is another case where

resonance occurs; there are two valid Lewis structures.



Self-Check EXERCISE 12.3 Ozone is a very important constituent of the atmosphere. At upper levels it

protects us by absorbing high-energy radiation from the sun. Near the earth’s

surface it produces harmful air pollution. Write the Lewis structure for

ozone, O3.

See Problems 12.63 through 12.68. ■

Now let’s consider a few more cases in Example 12.4.



EXAMPLE 12.4



Writing Lewis Structures: Summary

Give the Lewis structure for each of the following:



You may wonder how to decide

which atom is the central atom

in molecules of binary compounds. In cases where there is

one atom of a given element

and several atoms of a second

element, the single atom is

almost always the central atom

of the molecule.



a. HF



e. CF4



b. N2



f. NOϩ



c. NH3



g. NO3Ϫ



d. CH4

SOLUTION

In each case we apply the three steps for writing Lewis structures. Recall that

lines are used to indicate shared electron pairs and that dots are used to indicate nonbonding pairs (lone pairs). The table on page 379 summarizes our

results.



Self-Check EXERCISE 12.4 Write the Lewis structures for the following molecules:

a. NF3



d. PH3



g. NH4ϩ



b. O2



e. H2S



h. ClO3Ϫ



c. CO



f. SO42Ϫ



i. SO2



See Problems 12.55 through 12.68. ■



12.7 Lewis Structures of Molecules with Multiple Bonds



Molecule

or lon



Total Valence

Electrons



Draw Single

Bonds



Calculate Number

of Electrons

Remaining



Use Remaining

Electrons to

Achieve Noble

Gas Configurations



379



Check

Atom



Electrons



a. HF



1ϩ7 ϭ8



H



F



8Ϫ2 ϭ6



H



F



H

F



2

8



b. N2



5 ϩ 5 ϭ 10



N



N



10 Ϫ 2 ϭ 8



N



N



N



8



c. NH3



5 ϩ 3(1) ϭ 8



H



H

N



2

8



H



H

C



2

8



F



F

C



8

8



O ]ϩ



N

O



8

8



N

O



8

8



N

O



8

8



N

O



8

8



d. CH4



e. CF4



N



H



4 ϩ 4(1) ϭ 8



H



H



H

H



8Ϫ8 ϭ0



H



5 ϩ 6Ϫ1 ϭ 10



H



F



F

F



32 Ϫ 8 ϭ 24



F



O



N



5 ϩ 3(6)ϩ1 ϭ 24



10 Ϫ 2 ϭ 8



[ N



Ϫ



O

24 Ϫ 6 ϭ 18



N

O



C

F



O

g. NO3Ϫ



C



H



F

f. NOϩ



N

H



C



F



8Ϫ6 ϭ2



H



C



H



4 ϩ 4(7) ϭ 32



H



O



N

O



O

Ϫ



O



NO3Ϫ shows

resonance



N

O



O

Ϫ



O

N

O



O



Remember, when writing Lewis structures, you don’t have to worry

about which electrons come from which atoms in a molecule. It is best to

think of a molecule as a new entity that uses all the available valence electrons from the various atoms to achieve the strongest possible bonds. Think

of the valence electrons as belonging to the molecule, rather than to the individual atoms. Simply distribute all the valence electrons so that noble gas

electron configurations are obtained for each atom, without regard to the

origin of each particular electron.



▲



Some Exceptions to the Octet Rule

The idea that covalent bonding can be predicted by achieving noble gas electron configurations for all atoms is a simple and very successful idea. The

rules we have used for Lewis structures describe correctly the bonding in



380 Chapter 12 Chemical Bonding

most molecules. However, with such a simple model, some exceptions are inevitable. Boron, for example, tends to form compounds in which the boron

atom has fewer than eight electrons around it—that is, it does not have a

complete octet. Boron trifluoride, BF3, a gas at normal temperatures and

pressures, reacts very energetically with molecules such as water and ammonia that have unshared electron pairs (lone pairs).

H



H

O



→⎯⎯

→⎯⎯



Lone

⎯⎯⎯→ N

pairs



H



H

H



The violent reactivity of BF3 with electron-rich molecules arises because the

boron atom is electron-deficient. The Lewis structure that seems most consistent with the properties of BF3 (twenty-four valence electrons) is

F

Donald Clegg



B

F



Figure 12.10

When liquid oxygen is poured

between the poles of a magnet,

it “sticks” until it boils away.

This shows that the O2 molecule

has unpaired electrons (is

paramagnetic).



F



Note that in this structure the boron atom has only six electrons around it.

The octet rule for boron could be satisfied by drawing a structure with a double bond between the boron and one of the fluorines. However, experiments

indicate that each BOF bond is a single bond in accordance with the above

Lewis structure. This structure is also consistent with the reactivity of BF3

with electron-rich molecules. For example, BF3 reacts vigorously with NH3 to

form H3NBF3.

H

N ϩ〉



H

H



H



F

F → H

F



F

N



H



〉



F

F



Note that in the product H3NBF3, which is very stable, boron has an octet of

electrons.

It is also characteristic of beryllium to form molecules where the beryllium atom is electron-deficient.

The compounds containing the elements carbon, nitrogen, oxygen,

and fluorine are accurately described by Lewis structures in the vast majority of cases. However, there are a few exceptions. One important example is

the oxygen molecule, O2. The following Lewis structure that satisfies the

octet rule can be drawn for O2 (see Self-Check Exercise 12.4).

O



Paramagnetic substances have

unpaired electrons and are

drawn toward the space

between a magnet’s poles.



O



However, this structure does not agree with the observed behavior of oxygen.

For example, the photos in Figure 12.10 show that when liquid oxygen is

poured between the poles of a strong magnet, it “sticks” there until it boils

away. This provides clear evidence that oxygen is paramagnetic—that is, it

contains unpaired electrons. However, the above Lewis structure shows only

pairs of electrons. That is, no unpaired electrons are shown. There is no simple Lewis structure that satisfactorily explains the paramagnetism of the O2

molecule.

Any molecule that contains an odd number of electrons does not conform to our rules for Lewis structures. For example, NO and NO2 have eleven

and seventeen valence electrons, respectively, and conventional Lewis structures cannot be drawn for these cases.

Even though there are exceptions, most molecules can be described by

Lewis structures in which all the atoms have noble gas electron configurations, and this is a very useful model for chemists.