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potassium bromide. As we saw in Chapter 8, this occurs because when an
ionic solid dissolves, the ions are freed to move independently and can conduct an electric current.
The common polyatomic ions, which are listed in Table 4.4, are all held
together by covalent bonds.
Focus Questions
Sections 12.4–12.5
1. Why do metals lose electrons to form ions? When does a metal stop
losing ions?
2. Why does oxygen form an O2 ion and not an O3 ion?
3. Write the electron configurations for the pairs of atoms given below.
Use them to predict the formula for an ionic compound formed from
these elements.
a. Mg, S
c. Cs, F
b. K, Cl
d. Ba, Br
4. Why is aluminum foam useful in making cars more fuel-efficient?
5. Why are cations smaller than their parent atoms? Why are anions
larger?
6. How do polyatomic anions differ from simple anions?
Lewis Structures
Objective: To learn to write Lewis structures.
B
CHEMISTRY
Remember that the electrons in
the highest principal energy
level of an atom are called the
valence electrons.
onding 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 K ion has the [Ar] electron configuration and the Br 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:
K
[ Br ]
Noble gas
configuration [Ar]
Noble gas
configuration [Kr]
No dots are shown on the K ion because it has lost its only valence electron (the 4s electron). The Br ion is shown with eight electrons because it
has a filled valence shell.
12.6 Lewis Structures
371
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Lewis structures show only
valence electrons.
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 1s 2 and can be
represented by the following Lewis structure:
He
[He] configuration
CHEMISTRY
Carbon, nitrogen, oxygen, and
fluorine almost always obey the
octet rule in stable molecules.
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 (2s 22p6) of the neon atom are
represented by the Lewis structure. The 1s 2 electrons are core electrons and are not shown.
G. N. Lewis in his lab.
372
Next we want to develop some general procedures for writing Lewis structures for molecules. Remember that Lewis structures involve only the valence
electrons on atoms, so before we proceed, we will review the relationship of
Chapter 12 Chemical Bonding
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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 ns 2np4).
Group 6
Group
6
O
2s22p4
S
3s23p4
Se
4s24p4
Te
5s25p4
Similarly, all Group 7 elements have seven valence electrons (valence configuration ns 2np5).
Group 7
Group
7
F
2s22p5
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.
12.6 Lewis Structures
373
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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 second-row 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.
Example 12.2
Writing Lewis Structures: Simple Molecules
Write the Lewis structure of the water molecule.
Solution
We will follow the steps listed above.
Step 1 Find the sum of the valence electrons for H2O:
1
6
→
→
1
→
CHEMISTRY
The number of valence
electrons in an atom is the
same as the group number
on the periodic table for
representative elements.
H
H
O
(Group 1)
(Group 1)
(Group 6)
8 valence electrons
Step 2 Using a pair of electrons per bond, we draw in the two OH bonds,
using a line to indicate each pair of bonding electrons.
HOH
Note that
HOH 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 4) remain to be
distributed. Each hydrogen is satisfied with two electrons (duet rule), but
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Chapter 12 Chemical Bonding