Study Guide Chapters 12 – 14
Key
1. Define: electronegativity, dipole,
dipole moment, Van der Waals Forces.
1. Define: electronegativity, dipole,
dipole moment, Van der Waals Forces.
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•
•
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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 it’s 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).
2. State the differences and similarities between ionic,
covalent, and metallic bonds (see the “Four
Types of Bonding” table in your notebook).
2. State the differences and similarities between ionic,
covalent, and metallic bonds (see the “Four
Types of Bonding” table in your notebook).
Type of
Bonding
1. Metallic
2. Ionic
3. Polar
Covalent
4. Nonpolar
Covalent
Explanation of
Bonding
Nondirectional
bonds involving
positive metal ions
with free electrons
moving throughout
the metal.
The transfer of
electron(s) from a
metal atom to a
nonmetal atom
which results in an
attractive force
between a positive
metal ion and a
negative nonmetal
ion.
Electrostatic
attraction involving
unequal sharing of
electron(s) in
overlapping orbitals
between different
types of nonmetal
atoms.
Electrostatic
attraction involving
equal sharing of
electron(s) in
overlapping orbitals
between the same
type of nonmetal
atoms.
Form and
EXAMPLES
Properties.
Crystals with high copper,
melting points,
bronze, steel
shiny, malleable,
conductors,
Crystals with high
melting points
which are brittle
and act as
insulators.
NaCl, CuO,
FeBr2, etc
Molecules with
low melting points
which if solid are
brittle and act as
insulators.
H2O, CO2,
sucrose
(C6H12O6),
etc.
Molecules with
low melting points
which if solid are
brittle and act as
insulators.
H2, N2, O2
3. Contrast the number of shared pairs, the number
of electrons, the strength, and the length within
single, double, and triple bonds.
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.
4. What are the differences between shared
pairs and unshared pairs?
4. What are the differences between shared
pairs and unshared pairs?
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•
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 x’s, ’s or o’s.
They belong only to the atom which they
are placed on.
4. What are the differences between
shared pairs and unshared pairs?
4. What are the differences between
shared pairs and unshared pairs?
5. How does electronegativity vary within the
groups and periods of the periodic table?
5. How does electronegativity vary within the
groups and periods of the periodic table?
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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).
6. How can we predict the type of bond formed
between atoms by using (a) a periodic table and
(b) a table of electronegativities.
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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•
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.
7. How do differences in electronegativities
influence bond strength.
7. How do differences in electronegativities
influence bond strength.
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The greater the electronegativity
difference between two atoms the
stronger the bond is between them.
8. Complete the table:
8. Complete the table:
Number of atoms attached
to the central atom
4
2
Number of unshared pairs
attached to the central atom
Molecular
Shape
0
1
Trigonal planer
2
2
Trigonal pyramidal
linear
9. Contrast the attractive forces within solids,
liquids, and gases at room temperature.
9. Contrast the attractive forces within solids,
liquids, and gases at room temperature.
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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.
10. How do the size and polarity of molecules
affect their Van der Waals forces.
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.
11. PBr3
11. PBr3
11. PBr3
trigonal
pyramidal
11. PBr3
trigonal
pyramidal
polar
11. NO2-
11. NO2-
11. NO2-
Bent
11. NO2-
Bent
polar
11. ClF2+
11. ClF2+
11. ClF2+
Bent
11. ClF2+
Bent
polar
11. FNO2
11. FNO2
11. FNO2
Trigonal planer
11. FNO2
Trigonal planer
polar
11. N3-
11. N3-
11. N3-
linear
11. N3-
Linear
nonpolar
11. CF4
11. CF4
11. CF4
Tetrahedral
11. CF4
Tetrahedral
nonpolar