Chapter 8 Notes
Bonding and Molecular Structure:
Valence e- and Bonding
Covalent
Ionic
Bond Energy & Length
Structure, Shape & Polarity of Compounds
What is a Bond?
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A force that holds atoms together.
Why?
We will look at it in terms of energy.
Bond energy the energy required to break
a bond.
Why are compounds formed?
Because it gives the system the lowest
energy.
Covalent compounds?
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The electrons in each atom are attracted to
the nucleus of the other.
The electrons repel each other,
The nuclei repel each other.
The reach a distance with the lowest
possible energy.
The distance between is the bond length.
Thus Hydrogen is Diatomic!
Bond Formation
Covalent Character
e
Why Isn’t Helium Diatomic?
. .
He2
E
.
He + He
Inter-nuclear Distance
.
F+F
F2
2p ____ ____ ___
2s ____
F
___ ____ ____ 2p
____ 2s
F
Ionic Bonding
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An atom with a low ionization energy
reacts with an atom with high electron
affinity.
The electron moves.
Opposite charges hold the atoms together.
Li
+
Cl
1s22s1
[Ne] 3s23p5
2s ___
3p _____ _____ ___
1s _____
3s _____
[Ne]
Li + Cl
2s ___
3P _____ _____ _____
1s _____
3s _____
[Ne]
LiCl
2s ___
3P _____ _____ _____
1s _____
3s _____
[Ne]
Electronegativity
The difference between ionic
and covalent bonds.
Describes the relative ability
of an atom within a molecule
to attract a shared pair of
electrons to itself.
Electronegativity
Pauling electronegativity
values, which are unitless, are the norm.
Electronegativity
Range from 0.7 to 4.0
Bond: A - B
DEN = | ENA - ENB |
Bond Character
“Ionic Bond” - Principally Ionic
Character
“Covalent Bond” - Principally
Covalent
Character
Determining Principal
Character of Bond
DEN
~0
covalent
ionic
1.7
~4
F-F
Non-polar
EN = 0
DEN = |3.0 - 3.5|
= 0.5
N-O
N
O
Slightly polar
Ca - O
Ca
D EN = |1.0 - 3.5|
= 2.5
O
Ionic Bond with some
covalent character
Electronegativity
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D is known for almost every element
Gives us relative electronegativities of all
elements.
Tends to increase left to right.
decreases as you go down a group.
Noble gases aren’t discussed.
Difference in electronegativity between
atoms tells us how polar.
Zero
Covalent
Intermediate
Polar
Covalent
Large
Ionic
Covalent Character
decreases
Ionic Character increases
Electronegativity Bond
difference
Type
Dipole Moments
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A molecule with a center of negative
charge and a center of positive charge is
dipolar (two poles),
or has a dipole moment.
Center of charge doesn’t have to be on an
atom.
Will line up in the presence of an electric
field.
How It is drawn
d+ d-
H-F
Which Molecules Have Them?
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Any two atom molecule with a polar bond.
With three or more atoms there are two
considerations.
There must be a polar bond.
Geometry can’t cancel it out.
Ionic Radii -- Cations
Ionic Radii -- Anions
Molecular Polarity
MgBr2
Mg - Br
Br
EN = |1.2 - 2.8| = 1.6
Mg
Br
Covalent BOND w/much ionic
character, BUT NON-POLAR molecule
Lewis
Structures
The most important
requirement for the
formation of a stable
compound is that the
atoms achieve noble
gas e configuration
Valence Shell Electron
Pair Repulsion Model
(VSEPR)
The structure around a
given atom is determined
principally by minimizing
electron-pair repulsions
VSEPR
Electron Bond
pairs Angles
2
180°
Underlying
Shape
Linear
3
120°
4
109.5°
Tetrahedral
5
90° &
120°
6
90°
Trigonal
Bipyramidal
Octagonal
Trigonal Planar
LEWIS STRUCTURES
1 : draw skeleton of species
2 : count e- in species
3 : subtract 2
e
for each bond
in skeleton
4 : distribute remaining e-
Distinguish Between
ELECTRONIC
Geometry
&
MOLECULAR
Geometry
CH4
Bond angle = 109.50
Electronic geometry: tetrahedral
Molecular geometry: tetrahedral
+
H3O
Bond angle ~ 1070
Electronic geometry: tetrahedral
Molecular geometry: trigonal pyramidal
H2O
Bond angle ~ 104.50
Electronic geometry: tetrahedral
Molecular geometry: bent
NH2
Bond angle ~ 104.50
Electronic geometry: tetrahedral
Molecular geometry: bent
“Octet Rule” holds for
connecting atoms, but
may not for the central
atom.
BaI2
Bond angle =1800
Electronic geometry: linear
Molecular geometry: linear
BF3
Bond angle =1200
Electronic geometry: trigonal planar
Molecular geometry: trigonal planar
PF5
Bond angle = 1200 & 900
Electronic geometry: trigonal bipyramidal
Molecular geometry: trigonal bipyramidal
SF4
Bond angle = 1200 & 900
Electronic geometry: trigonal bipyramidal
Molecular geometry: see-saw
ICl3
Bond angle <= 900
Electronic geometry: trigonal bipyramidal
Molecular geometry: T-shape
I3
Bond angle = 1800
Electronic geometry: trigonal bipyramid
Molecular geometry: linear
PCl6
Bond angle = 900
Electronic geometry: octahedral
Molecular geometry: octahedral
BrF5
Bond angle ~ 900
Electronic geometry: octahedral
Molecular geometry: square pyramidal
ICl4
Bond angle = 900
Electronic geometry: octahedral
Molecular geometry: square planar
Actual shape
NonElectron Bonding Bonding
Pairs
Pairs
Pairs Shape
2
3
3
4
4
4
2
3
2
4
3
2
0
0
1
0
1
2
linear
trigonal planar
bent
tetrahedral
trigonal pyramidal
bent
Actual Shape
NonElectron Bonding Bonding
Pairs
Pairs
Pairs Shape
5
5
5
5
5
4
3
2
0
1
2
3
trigonal bipyrimidal
See-saw
T-shaped
linear
Actual Shape
NonElectron Bonding Bonding
Pairs
Pairs
Pairs Shape
6
6
6
6
6
6
5
4
3
2
0
1
2
3
1
Octahedral
Square Pyramidal
Square Planar
T-shaped
linear
What happens
when there are not
enough electrons
to “satisfy” the
central atom?
EXAMPLES
Ethene
Acetic Acid
Oxygen
Nitrogen
FAVORED LEWIS
STRUCTURES
1. formal charges closest to
zero
2. negative formal charge
is on the most
electronegative atom
EXAMPLES
Carbon dioxide
Thiocyanate ion
Sulfate ion
BOND
ENERGY &
LENGTH
Bond Energies
E = (Bonds Broken) – (Bonds Made)
Bonds form between atoms
because bonded atoms
exhibit a lower energy.
Thus, energy is required to
break bonds and energy is
released when bonds are
formed.
Bond Order = # bonds to a specific set
of elements
C-C the BO=1
C=C the BO=2
C
C the BO=3
Fractions are possible
COVALENT BONDS
Bond Dissociation
Energy
Bond Energy
(kJ/mol)
H-F 565
H-Cl 432
H-Br 366
H-I
299
Bond Energy
(kJ/mol)
Cl-Cl 242
Br-Br 193
I-I
151
Bond Energy
(kJ/mol)
CC
CC
CC
346
610
835
Bond Energy
(kJ/mol)
N  N 163
N  N 418
N  N 945
Energy
0
Internuclear Distance
Energy
0
Internuclear Distance
Energy
0
Internuclear Distance
Energy
0
Internuclear Distance
Energy
0
Bond Length
Internuclear Distance
Energy
Bond Energy
0
Internuclear Distance
Bond:
CC
CC
CC
Energy Length
(kJ/mol) (pm)
346
610
835
154
134
121
Bond:
CO
CO
Energy Length
(kJ/mol) (pm)
358
745
143
122