Covalent Bonding
• Sharing of electrons between two atoms
• Both atoms are usually nonmetals
• Two electrons are shared it is a single bond
• Four e - are shared it is a double bond
• Six e - are shared it is a triple bond
• Coordinate covalent bond is when both
electrons being shared come from one of the
atoms
• O 2
Lewis structures
Lewis structures
• O 2
• These two electrons are shared, so how many electrons
does each atom have
Lewis structures
• O 2
• Since each atom only has 7 e- some electrons have to
move to share more electrons until each has 8 electrons
Lewis structures
• O 2
• Since each atom only has 7 e- some electrons have to
move to share more electrons until each has 8 electrons
Lewis structures
• O 2
• Now each atom has 8 electrons so oxygen has a double
bond
How do you know where to put
atoms in a lewis structure
• The atom that has the most bonding
potential is the one that goes in the middle.
• This is the atom that has the most unpaired
electrons
• Atoms that have only one unpaired electron
must be on the outside of the structure.
ie) H,Li,Na,Cl,F
H 2 O
• Oxygen has two unpaired e- and hydrogen
only has one so which one goes in the
middle
• Oxygen
H 2 O
• Oxygen has two unpaired e- and hydrogen
only has one so which one goes in the
middle
• Oxygen
H 2 O
• Oxygen has two unpaired e- and hydrogen only
has one so which one goes in the middle
• Oxygen
• Hydrogen only has to have two electrons because
it is not big enough to hold 8
NH 3
• Nitrogen is in the middle because it has 3
unpaired electrons
NH 3
• Nitrogen is in the middle because it has 3
unpaired electrons
CO 2
• Carbon is in the middle because carbon has
4 unpaired electrons and oxygen only has
two
CO 2
• Carbon is in the middle because carbon has
4 unpaired electrons and oxygen only has
two
CO 2
• Carbon only has six electrons and each
oxygen now has seven so electrons have to
be moved to share
CO 2
• Carbon only has six electrons and each
oxygen now has seven so electrons have to
be moved to share
CO 2
• So CO 2 has double bonds
Coordinate Covalent Bond
• NH 4
+
•
Where does the last hydrogen have to go
Coordinate Covalent Bond
• NH 4
+
Since we are dealing with NH4+ what we
need to add is an H+, which has no e-. It
will share with the two e- on the top of the
Nitrogen.
Coordinate Covalent Bond
• NH 4
+
Since we are dealing with NH4+ what we
need to add is an H+, which has no e-. It
will share with the two e- on the top of the
Nitrogen.
Coordinate Covalent Bond
• NH 4
+
Since both electrons came from the nitrogen,
this is a coordinate covalent bond
VSEPR
• Valence Shell Electron Pair Repulsion
• Non bonding electrons around a central
atom will cause the molecule to bend
• Basic structures are tetrahedron, pyramidal,
bent, linear, (trigonal planar, square planar,
trigonal bipyramidal, octahedral, see saw.)
• H 2 O
VSEPR Examples
VSEPR Examples
• H 2 O
• So the molecule would look like this
• NH 3
VSEPR Examples
• NH 3
VSEPR Examples
• NH 3
VSEPR Examples
• CO 2
VSEPR Examples
• CO 2
VSEPR Examples
• CH 4
VSEPR Examples
Polarity
• Polarity is a difference in electronegativity
between two atoms that causes electrons to
not be shared equally.
• This causes one part of the molecule to
carry a slight positive charge and one side to
carry a slight negative charge
Polarity
• If the difference is
– 0.0-0.4 nonpolar covalent (shared equally)
Polarity
• If the difference is
– 0.0-0.4 nonpolar covalent (shared equally)
– 0.4-1.0 moderately polar covalent
Polarity
• If the difference is
– 0.0-0.4 nonpolar covalent (shared equally)
– 0.4-1.0 moderately polar covalent
– 1.0-2.0 polar covalent
Polarity
• If the difference is
– 0.0-0.4 nonpolar covalent (shared equally)
– 0.4-1.0 moderately polar covalent
– 1.0-2.0 polar covalent
– >2.0 ionic
Polarity
• The electrons will be around the atom with
the larger electronegativity more often
because larger electronegativity means that
the atom wants the electrons more.
• This means that the atom with the larger
electronegativity will have a partial negative
charge
Polarity
• The electrons will be around the atom with
the larger electronegativity more often
because larger electronegativity means that
the atom wants the electrons more.
• This means that the atom with the larger
electronegativity will have a partial negative
charge
Polarity
• The electrons will be around the atom with the
larger electronegativity more often because larger
electronegativity means that the atom wants the
electrons more.
• This means that the atom with the larger
electronegativity will have a partial negative
charge
• The other atom will have a partial positive charge
• H 2 O
Polarity Examples
• H 2 O
Polarity Examples
• H 2 O
Polarity Examples
• H 2 O
Polarity Examples
• CO 2
Polarity Examples
Polarity of Molecule
• If a molecule has partial positives and
partial negatives the molecule may be polar,
has a positive side and a negative side. If
you can separate all the + from the – by one
plane or one line it is polar.
• Shape of the molecule will make a
difference as to whether it is or is not polar
Polarity of Molecule
Intermolecular Attractions
• Attraction between two molecules (weak
bonds between two molecules)
• Van der Waals forces are the weakest
attractions and include dispersion forces and
dipole interactions.
Intermolecular Attractions
• Van der Waals forces are the weakest
attractions and include dispersion forces and
dipole interactions.
– Dispersion forces are caused by motion of
electrons creating very small electrical charges
Intermolecular Attractions
• Van der Waals forces are the weakest
attractions and include dispersion forces and
dipole interactions.
– Dispersion forces are caused by motion of
electrons creating very small electrical charges
– Dipole interactions are when the partial positive
of one molecule interacts with the partial
negative of another
Intermolecular Attractions
• Hydrogen bonding is a stronger bond than
van der Waals.
• It is an attractive force between hydrogen of
one molecule and the unpaired electrons on
a highly electronegative atom of another
molecule.
• Between hydrogen and either F, O, or N
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