Study Guide Chapters 12

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Study Guide Chapters 12 14

• Key

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1. Define electronegativity, dipole, dipole
moment, Van der Waals Forces.
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1. Define electronegativity, dipole, dipole
moment, Van der Waals Forces.

• electronegativity The tendency of a bonded atom
  to attract electrons towards itself. (example
  when F bonds to O the F pulls the electrons
  closer to it because its electronegativity is
  higher.)
• Dipole a polar molecule
• Dipole moment measurement of the amount of
  polarity. (example a molecule that is more
  polar would have a greater dipole moment).
• Van der Waals Forces Attractive forces between
  adjacent molecules. (example the bigger a
  molecule is and the more polar it is the better
  it is able to attract adjacent molecules).

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2. State the differences and similarities
between ionic, covalent, and metallic bonds (see
the Four Types of Bonding table in your
notebook).
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2. State the differences and similarities
between ionic, covalent, and metallic bonds (see
the Four Types of Bonding table in your
notebook).
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3. Contrast the number of shared pairs, the
number of electrons, the strength, and the length
within single, double, and triple bonds.

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3. Contrast the number of shared pairs, the
number of electrons, the strength, and the length
within single, double, and triple bonds.

• Single bonds have one shared pair of electrons
  (two shared electrons). They are the weakest and
  longest of the covalent bonds. Triple bonds have
  three shared pairs of electrons (6 shared
  electrons). They are the strongest and shortest
  of the covalent bonds. Double bonds have two
  shared pairs of electrons (four shared
  electrons). They have a strength and length
  between that of single and triple bonds.

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4. What are the differences between shared pairs
and unshared pairs?
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4. What are the differences between shared pairs
and unshared pairs?

• Shared pairs of electrons are represented in
  Lewis structures by s. They represent bonds
  and belong to both atoms which they connect.
• Unshared pairs (lone pairs) also called lone
  pairs are represented in Lewis structures by a
  pair of xs, ?s or os. They belong only to
  the atom which they are placed on.

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4. What are the differences between shared pairs
and unshared pairs?
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4. What are the differences between shared pairs
and unshared pairs?
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5. How does electronegativity vary within the
groups and periods of the periodic table?
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5. How does electronegativity vary within the
groups and periods of the periodic table?

• The closer an atom is to F in the periodic
  table the higher the electronegativity.
  Therefore electronegativities increase as we move
  up and to the right in the periodic table.
  (Remember that the noble gases have no
  electronegativities).

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6. How can we predict the type of bond formed
between atoms by using (a) a periodic table and
(b) a table of electronegativities.
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6. How can we predict the type of bond formed
between atoms by using (a) a periodic table and
(b) a table of electronegativities.

• A bond between a metal and a nonmetal is ionic.
  A bond between metals is metallic. A bond
  between different nonmetal atoms is polar
  covalent. A bond between the same nonmetal atoms
  is nonpolar covalent. We can also determine
  double and triple bonds from the C, N, and O
  groups of the periodic table.
• The electronegativity difference can be used to
  determine bond type as well.
• If the electronegativity difference is greater
  than 1.7 between two atoms the bond between them
  is ionic.
• If the electronegativity difference is less than
  0.3 the bond is nonpolar covalent.
• If the electronegativity is between 0.3 and 1.7
  the bond is polar covalent.

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7. How do differences in electronegativities
influence bond strength.
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7. How do differences in electronegativities
influence bond strength.

• The greater the electronegativity difference
  between two atoms the stronger the bond is
  between them.

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8. Complete the table
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8. Complete the table
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9. Contrast the attractive forces within solids,
liquids, and gases at room temperature.
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9. Contrast the attractive forces within solids,
liquids, and gases at room temperature.

• At any given temperature a solid has the greatest
  attractive forces and a gas has the least. The
  attractive forces in a liquid are somewhere in
  between.

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10. How do the size and polarity of molecules
affect their Van der Waals forces.
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10. How do the size and polarity of molecules
affect their Van der Waals forces.

• The larger and more polar a molecule is the
  greater its Van der Waals forces.

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11. PBr3
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11. PBr3
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11. PBr3
trigonal pyramidal
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11. PBr3
trigonal pyramidal polar
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11. NO2-
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11. NO2-
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11. NO2-

• Bent

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11. NO2-

• Bent
• polar

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11. ClF2
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11. ClF2
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11. ClF2

• Bent

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11. ClF2

• Bent polar

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11. FNO2
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11. FNO2
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11. FNO2

• Trigonal planer

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11. FNO2

• Trigonal planer polar

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11. N3-
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11. N3-
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11. N3-

• linear

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11. N3-

• Linear
• nonpolar

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11. CF4
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11. CF4
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11. CF4

• Tetrahedral

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11. CF4

• Tetrahedral nonpolar