Student Interactive PPT

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Chemistry 6/e
Steven S. Zumdahl
and Susan A. Zumdahl
Chapter 8:
BONDING:
GENERAL
CONCEPTS
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1
Bonding: General Concepts
• Types of Chemical
Bonds
• React 1
• Electronegativity
• React 2 • 3
• Bond Polarity and
Dipole Moments
• React 4 • 5 • 6
• Ions: Electron
Configurations and
Sizes
• React 7
• Formation of Binary
Ionic Compounds
• React 8 • 9 • 10 • 11
• React 12 • 13 • 14
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2
Questions to Consider
What is meant by the term “chemical
bond?”
Why do atoms bond with each other to
form molecules?
How do atoms bond with each other to
form molecules?
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3
Types of Chemical Bonds
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4
The Interaction of
Two Hydrogen
Atoms
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5
Energy profile as a function of the distance
between the nuclei of the hydrogen atoms
As the atoms approach each other (right side of graph),
the energy decreases until the distance reaches 0.074
nm (0.74 Å) and then begins to increase again due to
repulsions.
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6
Key Ideas in Bonding
Ionic Bonding: Electrons are transferred
Covalent Bonding: Electrons are shared
equally
What about intermediate cases?
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7
The Effect of an Electric Field on Hydrogen
Fluoride Molecules
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8
What is meant by the term “chemical
bond”?
Why do atoms bond with each other to
form molecules?
How do atoms bond with each other to
form molecules?
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9
Electronegativity
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10
If lithium and fluorine react, which has
more attraction for an electron? Why?
In a bond between fluorine and iodine,
which has more attraction for an electron?
Why?
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11
What is the general trend for
electronegativity across rows and down
columns on the periodic table?
Understand this trend, do not merely
memorize it.
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12
The Pauling Elecronegativity Values
Electronegativity generally increases across a
period and decreases down a group.
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13
The Relationship Between
Electronegativity and Bond Type
Electronegativity
Difference in the
Bonding Atoms
Zero
Covalent
Intermediate
Polar Covalent
Large
Ionic
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Covalent character
decreases; ionic
character increases
Bond
Type
14
Bond Polarity and
Dipole Moments
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15
Arrange the following bonds from most to
least polar:
I.
II.
III.
N-F
C-F
H-Cl
O-F
N-O
B-Cl
C-F
Si-F
S-Cl
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16
Which of the following bonds would be the
least polar yet still be considered polar
covalent?
Mg-O C-O
O-O
Si-O
N-O
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17
Which of the following bonds would be the
most polar without being considered
ionic?
Mg-O C-O
O-O
Si-O
N-O
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18
Ions: Electron Configurations
and Sizes
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19
Choose an alkali metal, an alkaline metal, a noble gas,
and a halogen so that they constitute an isoelectronic
series when the metals and halogen are written as their
most stable ions.
• What is the electron configuration for each species?
• Determine the number of electrons for each species.
• Determine the number of protons for each species.
• Rank the species according to increasing radius.
• Rank the species according to increasing ionization
energy.
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20
Sizes of Ions
Related to Positions
of the Elements in
the Periodic Table
Note that size generally
increases down a group.
Also note that in a series
of isoelectronic ions, size
decreases with
increasing atomic
number. The ionic radii
are given in units of
picometers.
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21
What we can “read” from the
periodic table:
Trends for
– Atomic size
– Ion radius
– Ionization energy
– Electronegativity
Electron configurations
Predicting formulas for ionic compounds
Ranking polarity of covalent bonds
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22
Formation of Binary Ionic
Compounds
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23
Comparison of
the energy
changes
involved in the
formation of
solid sodium
fluoride and
solid
magnesium
oxide.
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24
The relationship between the ionic character
of a covalent bond and the electronegativity
difference of the bonded atoms
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25
For a given molecule we want to:
• Describe the valence electron
arrangement in the molecule (Lewis
structure)
• Predict the geometry (VSEPR)
• Specify the atomic orbitals involved in
the molecule
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26
Lewis Structures
1. Sum the valence electrons.
2. Place bonding electrons between pairs
of atoms.
3. Atoms usually have noble gas
configurations.
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27
Draw a Lewis structure for each of the
following molecules:
H2
N2
O2
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F2
28
Draw a Lewis structure for each of the
following molecules:
H 2O
NH3
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29
Draw a Lewis structure for each of the
following molecules:
CO
CO2
H3COH
BF3
C 2H 6O
PCl5
NO3SF6
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30
VSEPR
To approximate molecular geometries:
1. Write Lewis structure.
2. Count electron pairs.
3. Choose an arrangement that
minimizes electron-pair repulsions.
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31
Determine the shape of each of the
following molecules including bond
angles:
HCN
PH3
SF4
O3
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KrF4
32
If you are asked to determine the shape
of a molecule, what is always the first
step?
How do we deal with multiple bonds with
VSEPR theory?
If more than one atom can break the octet
rule in a molecule, what do we do?
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33
True or false:
A molecule that has polar bonds will
always be polar.
If true, explain why.
If false, provide a counterexample.
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34
True or false:
Lone pairs make a molecule polar.
If true, explain why.
If false, provide a counterexample.
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35
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