Chapter 8(a)
Bonding:
General Concepts
Quartz grows in beautiful, regular crystals.
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8a–2
Figure 8.1: (a) The interaction of two hydrogen
atoms. (b) Energy profile as a function of the
distance between the nuclei of the hydrogen
atoms.
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8a–3
Figure 8.2: The effect of an electric field on hydrogen fluoride molecules. (a)
When no electric field is present, the molecules are randomly oriented.
(b) When the field is turned on, the molecules tend to line up with their
negative ends toward the positive pole and their
positive ends toward the negative pole.
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8a–4
Figure 8.3: The Pauling electronegativity values.
Electronegativity generally increases across a
period and decreases down a group.
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8a–5
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8a–6
Figure 8.4: (a) The charge distribution in the
water molecule. (b) The water molecule in an
electric field.
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8a–7
Figure 8.5: (a) The structure and charge distribution of the
ammonia molecule. The polarity of the N—H bonds occurs
because nitrogen has a greater electronegativity than
hydrogen. (b) The dipole moment of the ammonia molecule
oriented in an electric field.
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8a–8
Figure 8.6: (a) The carbon dioxide molecule.
(b) The opposed bond polarities cancel out, and
the carbon dioxide has no dipole moment.
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8a–9
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8a–10
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8a–11
Figure 8.7: Sizes of ions related to positions
of the elements on the periodic table.
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8a–12
Chapter 8(b)
Bonding:
General Concepts
(cont’d)
Lithium fluoride
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8a–14
Figure 8.8: The energy changes involved in
the formation of solid lithium fluoride from its
elements.
Figure 8.9:
The structure
of lithium
fluoride.
Figure 8.10:
Comparison
of the energy
changes
involved in the
formation of
solid sodium
fluoride and
solid
magnesium
oxide.
Figure 8.11: The
three possible
types of bonds
(a) a covalent
bond formed
between identical
atoms (b) a polar
covalent bond,
with both ionic and
covalent
components; and
(c) an ionic bond
with no electron
sharing.
Figure 8.12: The relationship between the ionic
character of a covalent bond and the
electronegativity difference of the bonded atoms.
Molten NaCl
conducts an
electric
current,
indicating the
presence of
moblie Na+
and Cl- ions.
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8a–21
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8a–22
Figure 8.14:
The molecular
structure of
methane. The
tetrahedral
arrangement of
electron pairs
produces a
tetrahedral
arrangement of
hydrogen
atoms.
Figure 8.15:(a) The tetrahedral arrangement of electron pairs around the
nitrogen atom in the ammonia molecule. (b) Three of the electron pairs
around nitrogen are shared with hydrogen atoms as shown and one is a
lone pair. Although the arrangement of electron pairs is tetrahedral, as in
the methane molecule, the hydrogen atoms in the ammonia molecule
occupy only three corners of the tetrahedron. A lone pair occupies the
fourth corner. (c) Note that molecular geometry is
trigonal pyramidal, not tetrahedral.
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8a–24
Chapter 8(c)
Bonding:
General Concepts
(cont’d)
Figure 8.16: (a) The tetrahedral arrangement of the four
electron pairs around oxygen in the water molecule. (b) Two
of the electron pairs are shared between oxygen and the
hydrogen atoms and two are lone pairs. (c) The V-shaped
molecular structure of the water molecule.
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8a–26
Figure 8.17: The bond angles in the CH4,
NH3, and H2O molecules.
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8a–27
Figure 8.18: (a) In
a bonding pair of
electrons, the
electrons are
shared by two
nuclei. (b) In a lone
pair, both electrons
must be close to a
single nucleus and
tend to take up
more of the space
around that atom.
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8a–29
-
Molecular structure of PCl6
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8a–30
Octahedral
electron
arrangement
for Xe
Figure 8.19: Possible electron-pair
arrangements for XeF4.
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8a–32
Figure 8.20: Three possible arrangements of
the electron pairs in the I3 ion.
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8a–33
Figure 8.21: The
molecular structure
of methanol.
(a) The arrangement
of electron pairs and
atoms around the
carbon atom.
(b) The arrangement
of bonding and lone
pairs around the
oxygen atom.
(c) The molecular
structure.