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370 Chapter 12 Chemical Bonding
12.6 Lewis Structures
OBJECTIVE:
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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.