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Lecture

CHEM 212 Lecture Notes - Electron Diffraction, Lone Pair, Bond Length


Department
Chemistry
Course Code
CHEM 212
Professor
Richard Oakley

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We now turn from the use of quantum mechanics and its description of the atom to an
elementary description of molecules. Although most of the discussion of bonding in this
book uses the molecular orbital approach to chemical bonding, simpler methods that
provide approximate pictures of the overall shapes and polarities of molecules are also
very useful. This chapter provides an overview of Lewis dot structures, valence shell
electron pair repulsion
(VSEPR),
and related topics. The molecular orbital descriptions
of some of the same molecules are presented in Chapter
5
and later chapters, but the
ideas of this chapter provide a starting point for that more modern treatment. General
chemistry texts include discussions of most of these topics; this chapter provides a re-
view for those who have not used them recently.
Ultimately, any description of bonding must be consistent with experimental data
on bond lengths, bond angles, and bond strengths. Angles and distances are most fre-
quently determined by diffraction (X-ray crystallography, electron diffraction, neutron
diffraction) or spectroscopic (microwave, infrared) methods. For many molecules, there
is general agreement on the bonding, although there are alternative ways to describe it.
For some others, there is considerable difference of opinion on the best way to describe
the bonding. In this chapter and Chapter
5,
we describe some useful qualitative ap-
proaches, including some of the opposing views.
I'
E
f
1:
3-1
Lewis electron-dot diagrams, although very much oversimplified, provide a good
[
LEWIS ELECTRON-
starting point for analyzing the bonding in molecules. Credit for their initial use goes
DOT DIAGRAMS
to G.
N.
~ewis,' an American chemist who contributed much to thermodynamics and
chemical bonding in the early years of the 20th century. In Lewis diagrams, bonds
between two atoms exist when they share one or more pairs of electrons. In addition,
some molecules have nonbonding pairs (also called lone pairs) of electrons on atoms.
'G.
N.
Lewis,
J.
Am.
Chem. Soc., 1916,38,762; Valence and the Structure of Atonzs and Molecules,
Chemical Catalogue Co.,
New
York,
1923.
5
1

Only pages 1-3 are available for preview. Some parts have been intentionally blurred.

52
Chapter
3
Simple Bonding Theory
These electrons contribute to the shape and reactivity of the molecule, but do not
directly bond the atoms together. Most Lewis structures are based on the concept that
eight
valence electrons
(corresponding to
s
and
p
electrons outside the noble gas core)
form a particularly stable arrangement, as in the noble gases with
sZp6
configurations.
An exception is hydrogen, which is stable with two valence electrons. Also, some
molecules require more than eight electrons around a given central atom.
A more detailed approach to electron-dot diagrams is presented in Appendix
D.
Simple molecules such as water follow the
octet rule,
in which eight electrons
surround the oxygen atom. The hydrogen atoms share two electrons each with the oxy-
gen, forming the familiar picture with two bonds and two lone pairs:
Shared electrons are considered to contribute to the electron requirements of both
atoms involved; thus, the electron pairs shared by
H
and
0
in the water molecule are
counted toward both the 8-electron requirement of oxygen and the 2-electron require-
ment of hydrogen.
Some bonds are double bonds, containing four electrons, or triple bonds, contain-
ing six electrons:
3-1
-1
RESONANCE
In many molecules, the choice of which atoms are connected by multiple bonds is arbi-
trary. When several choices exist, all of them should be drawn. For example, as shown
in Figure 3-1, three drawings (resonance structures) of
~0-3~-
are needed to show the
double bond in each of the three possible C-0 positions. In fact, experimental evi-
dence shows that all the C-0 bonds are identical, with bond lengths (129 pm) be-
tween double-bond and single-bond distances
(1
16
pm and 143 pm respectively); none
of the drawings alone is adequate to describe the molecular structure, which is a combi-
nation of all three, not an equilibriu~n between them. This
is
called
resonance
to signi-
fy that there is more than one possible way in which the valence electrons can be placed
in a Lewis structure. Note that in resonance structures, such as those shown for ~03~-
in Figure 3-1, the electrons are drawn in different places but the atomic nuclei remain in
fixed positions.
The species
CO~~-,
NOs-,
and SO3, are
isoelectronic
(have the same electronic
structure). Their Lewis diagrams are identical, except for the identity of the central atom.
When a molecule has several resonance structures, its overall electronic energy is
lowered, making it more stable. Just as the energy levels of a particle in a box are low-
ered by making the box larger, the electronic energy levels of the bonding electrons are
lowered when the electrons can occupy a larger space. The molecular orbital descrip-
tion of this effect is presented in Chapter
5.
:G:
7
2-
:G:
72-
.
.
d
7
.
.
d
.
.
7
FIGURE
3-1
Lewis Diagrams for
.
.
*.c
-
.
.C..
.. ..
c:.
c01>-.
-9.
-.@
0
-
'
.o:
..0.
'
0

Only pages 1-3 are available for preview. Some parts have been intentionally blurred.

52
Chapter
3
Simple
Bonding Theory
These electrons contribute to the shape and reactivity of the molecule, but do not
directly bond the atoms together. Most Lewis structures are based on the concept that
eight
valence electrons
(corresponding to
s
and
p
electrons outside the noble gas core)
form a particularly stable arrangement, as in the noble gases with
s2p6
configurations.
An exception is hydrogen, which is stable with two valence electrons. Also, some
molecules require more than eight electrons around a given central atom.
A
more detailed approach to electron-dot diagrams is presented in Appendix D.
Simple molecules such as water follow the
octet rule,
in which eight electrons
surround the oxygen atom. The hydrogen atoms share two electrons each with the oxy-
gen, forming the familiar picture with two bonds and two lone pairs:
Shared electrons are considered to contribute to the electron requirements of both
atoms involved; thus, the electron pairs shared by
H
and
0
in the water molecule are
counted toward both the 8-electron requirement of oxygen and the 2-electron require-
ment of hydrogen.
Some bonds are double bonds, containing four electrons, or triple bonds, contain-
ing six electrons:
3-1-1
RESONANCE
In many molecules, the choice of which atoms are connected by multiple bonds is arbi-
trary. When several choices exist, all of them should be drawn. For example, as shown
in Figure
3-1,
three drawings (resonance structures) of
~0~~-
are needed to show the
double bond in each of the three possible C-0 positions. In fact, experimental evi-
dence shows that all the C-0 bonds are identical, with bond lengths (129 pm) be-
tween double-bond and single-bond distances
(1
16 pm and 143 pm respectively); none
of the drawings alone is adequate to describe the molecular structure, which is a combi-
nation of all three, not an equilibrium between them. This is called
resonance
to signi-
fy that there is more than one possible way in which the valence electrons can be placed
in a Lewis structure. Note that in resonance structures, such as those shown for ~03~-
in Figure
3-1,
the electrons are drawn in different places but the atomic nuclei remain in
fixed positions.
The species
CO~~-,
NO3-, and SO3, are
isoelectronic
(have the same electronic
structure). Their Lewis diagrams are identical, except for the identity of the central atom.
When a molecule has several resonance structures, its overall electronic energy is
lowered, making it more stable. Just as the energy levels of a particle in a box are low-
ered by making the box larger, the electronic energy levels of the bonding electrons are
lowered when the electrons can occupy a larger space. The molecular orbital descrip-
tion of this effect is presented in Chapter
5.
:a:
72-
.!.
72-
:a:
72-
. .
+
7
.
.
+
.
.
7
FIGURE
3-1
Lewis Diagrams for
:O'.'
.c.
,
.
..c..
..
..
c:.
c01~-.
.
.
.g.
0:
'
.0.:
e.0.
'
Q:
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