CHEMICAL BONDING
Hydrogen Bonding
Chemical Bonding Unit Skills
• Compare ionic, covalent, and metallic
bonds
• Draw molecular structures
• Distinguish between polar and nonpolar
compounds
• Explain the affect of hydrogen bonds on
compounds and how the are formed
This unit corresponds with chapter 12
Chemical Bonds
• Force of attraction between two atoms
• Several kinds of bonds:
– Ionic bonds
• Network solids
– Covalent bonds
• Macromolecules
• Covalent I & II
– Metallic bonds
+
metal
Ionic Bonds
nonmetal
• Occur in ionic compounds generally
between a metal and a nonmetal
• Formed when electrons are transferred
• Strong bond
• High melting point
• Generally solids
• Poor conductors of electricity as a solid,
but good conductors when dissolved
Covalent Bonds
• Generally occur in molecular compounds
between two nonmetals
• Formed when electrons are shared
• Weak bonds ***
• Low melting points
• Usually liquids or gases
• Poor conductors of electricity
Polar covalent bonding
Occurs when 2 slightly different atoms share
electrons unequally to be more stable.
The electrons are not completely transferred
but an unequal sharing results. We use these
symbols to show which atom has a stronger
attraction for the electrons.
Which of the following compounds
contains ionic bonds?
1.
2.
3.
4.
NaCl
NO
F2
SO3
Sharing in covalent bonds
• Electronegativity determines how electrons are
shared
• Uneven sharing = polar bond
• Even sharing = nonpolar bond
•••
•
Electronegativity
What
element the
hasattraction,
the
highest
Theelements
stronger
theelectronegativity?
higher
the number.
The
have
been
assigned
numbers
according to their
What
the lowest?
ability has
to “attract”
electrons to itself from the shared bond.
A bond’s polarity depends on the
difference in the electronegativity values
of the atoms in the bond.
HF
2.1
4.0
H Cl
2.1
3.0
4.0 The HF
-2.1 bond
is
1.9
3.0
-2.1
0.9
more
polar
Metallic Bonds
• Occur between atoms of a metal of the
same element.
• Strong bonds
• Have all of the properties of metals
discussed with the periodic table
Metals form a regular
array of atoms. But the
valence electrons form
a “sea” of electrons also
called “delocalized”.
Intermolecular Attractions
Why is carbon a solid at room
temperature and water a liquid?
Dipoledipole
attractions
are 1% as
strong as
regular
bonds.
Different
attractions
between
molecules
determine
the
properties
and states
of matter.
Hydrogen Bonds
• Bond that occurs
between molecules
containing hydrogen
and an atom with a
high electronegativity
(usually N, O, F, Cl, or
S)
Hydrogen bonds are strong
intermolecular attractions…
about 10 x stronger than
dipole-dipole attractions
London dispersion
forces are another
Even non-polar
matter.
molecules can form
weak attractions. This
is due to the uneven
distributions of electrons
at any one time.
Lewis Structures (Dot Diagrams)
• Used to show the arrangement of atoms in
molecules or ions
• Follow the steps to draw Lewis structures
• But first a quick review
Let’s try bromine.
Notice: bromine now has 8 electrons!
2
10
5
4s 3d 4p
Br [Ar]
2
10
6
Br [Ar] 4s 3d 4p
How many valence electrons?
What would the bromine ion look like?
H2
H H
By sharing electrons, each hydrogen has, in
effect, 2 electrons or a filled valence shell
(stable duet)
Why doesn’t Helium form bonds?
He
How about F2?
F F
Each fluorine has 7 valence electrons.
That is 1 bonding pair and 3 “lone pairs” or
“unshared pairs”
Does Neon form bonds?
Ne
Now let’s look at families
on the periodic table.
Li
Na
K
Rb
Cs
Fr
O
[He] 2s1
S
[Ne] 3s1
[Ar] 4s1
[Kr] 5s1
Se
[Xe] 6s1
[Rn] 7s1
Te
H O H
H - O- H
Count the total valence electrons: 1 + 6 + 1 = 8
Place the electrons to form a bond between
atoms and to complete the duet and octet.
Often, instead of showing 2 dots between
atoms for a bond, we draw a bar.
Sometimes we have multiple bonds!
CO2
O=C=O
O-C-O
Count the total valence electrons:6 + 4 + 6 = 16
The counts work but both oxygen atoms do not
have a stable octet.
This is the best arrangement for CO2,
but is there another?
O=C=O
O-C=O
O=C−O
These are called
resonance structures. A
molecule shows
resonance when more
than one Lewis structure
can be drawn. We will
discuss later which one
is the most stable.
Try another:
CN
C-N
Count the total valence electrons:4 + 5 + 1= 10
Are there any resonance structures?
Try these:
HF
•• •
H-F
•
••
N2
•• N = N••
••
NH3
CH4
H–N–H
H
H
H–C–H
H
| |
F
| F – C – F|
4 | F|
CF
+
NO •• N
-
= O••
-NO3| O – N = O|
- | | O
-
More to practice on your own!
a. NF3
b. O2
c. CO
d. PH3
e. H2S
f. SO42g. NH4+
h. ClO3i. SO2
Steps for drawing Lewis structures
1. Count the number of valence electrons in
the molecule or ion
2. Arrange the atoms around a central atom
3. Put a pair of electrons (2 dots) where
each bond occurs
4. Put the remaining electrons around each
atom so all have 8 except hydrogen
which can only have 2
Some exceptions to the octet rule
• A few atoms seem to be stable being
“electron deficient”
• The atoms are Be, B, Mg and Al
• So we can have molecules like BF3
F
I’ve left off the electrons
around the fluorine atoms
due to time & space.
F
B
F
VSEPR Model
•
•
•
•
•
Valence Shell Electron Pair Repulsion
When bonds form, several forces come in to
play…
The bonding electrons let atoms achieve a
noble gas configuration (lowering energy).
A small concentration of negative charge forms
between the two atoms.
Also, there may be unbonded pairs of electrons.
These create a very large concentration of
negative charge and these repel other areas of
negative charge.
Basic structures & predicting
•
•
•
•
•
•
•
•
•
•
•
Linear
Trigonal planar
Bent & v-shape
Tetrahedral
Trigonal pyramidal
Trigonal bipyramidal
See saw
T-shape
Octahedral
Square pyramidal
Square planar
We’ll take
these one
at a time.
Notation
AXE
• A represents the central atom
• X represents the ligands (atoms
bonded to central atom)
• E represents the unbonded
electron pairs
Linear
Cl Be Cl
Bond
angle
180o
BeCl2
Both electron
Hybridization:
and
2 0
sp
molecular
geometry are central atom
ligands
linear.
unbonded electron pairs
AX E
Trigonal planar
F F
B
F
Both electron
and molecular
geometry are
trigonal planar.
A X3 E0
Bond
angle
120o
BF3
Hybridization:
sp2
central atom
ligands
unbonded electron pairs
Bent
O
S
Electron
geometry is
trigonal planar
and molecular
geometry is bent
or v-shape
O
Bond angle
less than
120o
SO2
A X2 E1
central atom
ligands
unbonded
electron pairs
Hybridization:
sp2
Tetrahedral
H
H CH
H
Both electron
and molecular
geometry is
tetrahedral.
Bond angle
109.5o
A X4 E0
CH4
Hybridization:
sp3
Trigonal Pyramidal
H NH
H
Electron geometry
is tetrahedral and
molecular
geometry is
trigonal pyramidal.
Bond angle
107o
A X3 E1
NH3
Hybridization:
sp3
Very Bent
H O
H
Electron geometry
is tetrahedral and
molecular
geometry is bent
or very bent.
Bond angle
104.5o
AX2 E2
H2O
Hybridization:
sp3
Trigonal bipyramidal
Leaving off the unshared electron pairs on F due to space.
F F
F P
F
F
Both the electron and
molecular geometry
is trigonal
bipyramidal
PF5
Bond angles
90o and 120o
AX5 E0
Hybridization:
sp3d
See saw
Leaving off the unshared electron pairs on F due to space.
F F
S
F
F
Electron geometry is
trigonal bipyramidal,
molecular geometry
is see saw or
distorted tetrahedral
SF4
Bond angles
90o and 120o
AX4 E1
Hybridization:
sp3d
T-shaped
Leaving off the unshared electron pairs on F due to space.
F F
ClF3
Bond angles
90o
Cl
F
Electron geometry is
trigonal bipyramidal,
molecular geometry
is T-shaped
AX3 E2
Hybridization:
sp3d
Linear
Leaving off the unshared electron pairs on F due to space.
F
XeF2
Bond angles
180o
Xe
F
Electron geometry is
trigonal bipyramidal,
molecular geometry
is linear
AX2 E3
Hybridization:
sp3d
Octahedral
Leaving off the unshared electron pairs on F due to space.
F
F
F
S
F F F
SF6
Bond angles
90o
Both electron
geometry and
molecular geometry
is octahedral
AX6 E0
Hybridization:
sp3d2
Square Pyramidal
Leaving off the unshared electron pairs on F due to space.
F
F
F
Br
F
F
BrF5
Bond angles
90o
Electron geometry is
octahedral,
molecular geometry
is square pyramidal
AX5 E1
Hybridization:
sp3d2
Square Planar
Leaving off the unshared electron pairs on F due to space.
F
F
Xe
F
F
XeF4
Bond angles
90o
Electron geometry is
octahedral,
molecular geometry
is square planar
AX4 E2
Hybridization:
sp3d2
Try these:
What is the molecular geometry of GeI4?
4 + 4(7) = 32
I
Tetrahedral
I
Ge I
I
AX4 E0
Try these:
What is the molecular geometry of GeI4?
4 + 4(7) = 32
I
Tetrahedral
I
Ge I
I
AX4 E0
Molecular Structure Worksheet
Molecules
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•
•
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Carbon tetrachloride
Phosphorus trichloride
Hydrogen sulfide
Chlorine molecule (Cl2)
Sulfur dichloride
Carbon monoxide
Ethane (C2H6)
Silicon dioxide
Hydrogen cyanide
Molecular Shapes
What happens when oil is put in
water?
•
Like Dissolves Like
Cell Membranes