14.2 Hybridization
(and quick review of Intermolecular
Forces-4.3)
Chemical Bonds
Intramolecular Forces
• Bonds within Molecules
– Ionic
– Covalent
– Metallic
Intermolecular Forces
• Bonds between Molecules
– van der Waals Forces
• London Forces
• Dispersion Forces
– Dipole-Dipole Forces
– Hydrogen Bonds
© 2009, Prentice-Hall, Inc.
Comparing Strengths
• Which are stronger-intermolecular forces or
intramolecular forces?
• What types of bonds are being broken when
melting an ionic compound? Covalent
Compound?
van der Waal’s Forces
aka
London Dispersion Forces
• Weakest intermolecular force
• Exist between all molecules
• A temporary dipole in one nonpolar molecule induces a
dipole in neighboring molecules—resulting in an
instantaneous attraction between the molecules.
© 2009, Prentice-Hall, Inc.
van der Waal’s Forces
aka
London Dispersion Forces
• Weakest intermolecular force
• Exist between all molecules
• A temporary dipole in one nonpolar molecule induces a
dipole in neighboring molecules—resulting in an
instantaneous attraction between the molecules.
© 2009, Prentice-Hall, Inc.
Dipole-Dipole Forces
• Polar molecules are
attracted to each other
– Electrostatic attractions
– The positive end of one is
attracted to the negative
end of the other and viceversa.
© 2009, Prentice-Hall, Inc.
Dipole-Dipole Forces
• The more polar the molecule, the higher its
boiling point.
© 2009, Prentice-Hall, Inc.
Hydrogen Bonding
• Interaction experienced when H is
bonded to highly electronegative
and small atoms like N, O, or F .
• Molecules are extremely polar
• Form unusually strong dipole-dipole
interaction
© 2009, Prentice-Hall, Inc.
Hydrogen Bonding
• Considerably stronger
than vDW or dipole-dipole
forces.
• Molecules that are
hydrogen bonded have
higher boiling and melting
points.
• Hydrogen bonded
molecules have greater
solubility in water.
© 2009, Prentice-Hall, Inc.
Warm-up 10/25
1. Determine the molecular shape of:
CH3CHCHCN
Now to 14.2-Hybridization
• Orbital Shapes
Sigma and Pi Bonds
– Sigma (σ) Bonds
End-to-end overlap
– Pi bonds (π) Bonds
Sideways overlap
Single, Double, and Triple Bonds
Explaining the bond strengths/lengths
• Single bond is longest/weakest
– Least area of overlap (just sigma)
• Double bond shorter/stronger
– 3 areas of overlap (one sigma and one pi)
– Pulls two nuclei closer together
• Triple bond shortest/strongest
– 5 areas of overlap (1 sigma, 2 pi)
– Two nuclei even closer together
Hybridization
• http://liakatas.org/chemblog/?page_id=17#Si
mulations
• CH4 - 1s2 2s2 2p2
– How can we explain the 4 equivalent bonds?
– Bonding electrons are in different orbitals
• 2 in the 2p orbitals and 2 in the 2s orbital
Hybrid orbitals are created
• Must be 4 equivalent orbitals
– The s and p orbitals involved in forming sigma
bonds (or containing lone pairs), mesh together
– Form equal number of highly directional hybrid
orbitals
• CH4 = four sp3 hybrid orbitals
CH4
• 4 sp3 orbitals
– sp3 orbitals look
like half p orbitals
Key Ideas
• When atoms join to form molecules (except H)
– Outer atomic orbitals produce hybrid orbitals
• Focus only on the central atom
• Results in same number of hybrid orbitals as
original orbitals involved
– Now they all have the same energy and are
arranged symmetrically
sp3 hybridization in water
In the case of water, the three 2p
orbitals of the oxygen atom are
combined with the 2s orbital to
form four sp3 hybrid orbitals. The
two non-bonded electron pairs
will occupy hybrid orbitals. Again
we need a hybrid orbital for each
atom and each pair of nonbonding electrons. Water has two
hydrogen atoms and two nonbonded pairs of electrons when
we draw the electron-dot formula.
sp2 hybridization in Boron Trifluoride
In the boron trifluoride molecule,
only three groups are arranged
around the central boron atom. In
this case, the 2s orbital is combined
with only two of the 2p orbitals (since
we only need three hybrid orbitals for
the three groups...thinking of groups
as atoms and non-bonding pairs)
forming three hybrid orbitals called
sp2 hybrid orbitals. The other porbital remains unhybridized and is at
right angles to the trigonal planar
arrangement of the hybrid orbitals.
The trigonal planar arrangement has
bond angles of 120o.
sp Hybridization
• BeCl2
Shapes to Hybrid Orbitals
Lewis Dot Structures to Hybrid Orbitals
• From Lewis Dot Structure:
– Number of orbitals around central atom = number
of charge centers
• Only care about sigma bonds so single, double and
triple bonds counted as one.
Example: H2O
-2 sigma bonds and 2 non-bonding pairs 
4 orbitals so sp3 hybridization
Homework
Alternative text Ch. 4
• Pg. 117
– #1, 2
• Pg. 118
– #1, 2