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Chapter 8(a)
Bonding: Bonding: General ConceptsGeneral Concepts
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Quartz grows in beautiful, regular crystals.
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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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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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Figure 8.3: The Pauling electronegativity values. Electronegativity generally increases across a
period and decreases down a group.
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Figure 8.4: (a) The charge distribution in the water molecule. (b) The water molecule in an
electric field.
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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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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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Figure 8.7: Sizes of ions related to positions of the elements on the periodic table.
Chapter 8(b)
Bonding: Bonding: General Concepts General Concepts
(cont’d)(cont’d)
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Lithium fluoride
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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Figure 8.14: The molecular structure of methane. The tetrahedral arrangement of electron pairs produces a tetrahedral arrangement of hydrogen atoms.
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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.
Chapter 8(c)
Bonding: Bonding: General Concepts General Concepts
(cont’d)(cont’d)
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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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Figure 8.17: The bond angles in the CH4, NH3, and H2O molecules.
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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Molecular structure of PCl6-
Octahedral electron arrangement for Xe
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Figure 8.19: Possible electron-pair arrangements for XeF4.
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Figure 8.20: Three possible arrangements of the electron pairs in the I3- ion.
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.