Molecular Polarity &
Intermolecular Forces
How to Predict Whether a
Molecule is Polar or Nonpolar
– You already learned how to determine
whether a BOND is polar or nonpolar using
a Table of Electronegativity.
– You learned to determine and draw the
correct molecular geometry of a molecule.
– You can put these two pieces together and
determine whether the overall molecule is
polar.
– Although you might think that a molecule
which has at least one polar bond would
automatically be a polar molecule, this
would be an incorrect assumption.
– There are lots of nonpolar molecules which
contain polar bonds. However, it is true that
a molecule with no polar bonds and no lone
pairs must be nonpolar.
– How is this possible? Because of the overall 3-D
shape!
– If the molecule and the polar bonds are
symmetrically arranged, they may cancel out.
– Thus, the molecule would be nonpolar with polar
bonds.
– Let’s take a look:
CCl4
H2 S
CH3Cl
SiH4
NH3
PH3
– So not only are the individual bond dipoles
important, but lone pairs of electrons affect the
molecular polarity as well.
– Lone pairs of electrons pull electron density away
from the central atom, so they have a dipole as well.
– We can draw an overall dipole for the entire
molecule, called the dipole moment.
– This dipole moment is the vector addition of the
individual bond dipoles as well as the pull of lone
pairs on electron density.
– If you look at a table of molecular geometries, there
are some generalizations about molecular polarity
which can be made:
• Your Turn: Determine whether CO2,
CF4 , CH2Cl2, and H2O are polar or
nonpolar
Partial Ionic Character
• We often say that a molecule like HF has
“partial” ionic character.
• Or we can say that an “ionic” compound
like AlN has “partial” covalent character.
• This can actually be calculated.
– If a molecule is polar, it has a nonzero dipole
moment. If a molecule is nonpolar, it has a
dipole moment of zero.
– The dipole moment is defined as:  = Qr
• where  is the Dipole moment in debye units (D),
Q is the charge in coulombs (C), and r is the
distance between the charges.
– The higher the dipole moment, the more
polar the molecule is or you could say that
the bonds have more ionic character.
– We can use this equation to calculate a
dipole moment, but as it is actually easy to
experimentally measure a molecule’s dipole
moment, we more often use the dipole
moment to calculate the partial charge or %ionic character of a molecule.
– HBr has a bond length of 141 pm and a
dipole moment of 0.79 D. Given that the full
charge of an electron is 1.60x10-19C and 1 D
= 3.336x10-30Cm, what is the %-ionic
character of the H-Br bond?
Intermolecular Forces
Intermolecular Forces
– You already know that CO2 is a gas at room
temperature, while water is a liquid, and
sucrose is a solid.
– Why? They are all molecular species.
– What holds water molecules together in the
liquid phase at room temperature?
– Or what determines what phase or state a
compound will exist in at room
temperature?
– Let’s review the 3 states of matter:
Gas
Liquid
Solid
no fixed volume
fixed volume
fixed volume
no fixed shape
no fixed shape
fixed shape, so
rigid
high densities
very low densities high densities
density varies
with T and P
not compressible
not compressible
rapid, random
motion
fluid, motion
little motion
high kinetic
motion
some kinetic
energy
little kinetic
energy
Intermolecular Forces
– So what determines the state of a
compound at room temp?
– The strength of the attractions
between separate molecules: the
stronger the attraction, the more likel
the compound will be a solid, the
weaker the attractions, the more
likely it will be a gas.
Intermolecular Forces
– These attractive forces “glue” solids
or liquids together.
– What are these forces?
Intermolecular Forces
– The forces that hold individual
molecules together in the solid or
liquid phases are called intermolecular
forces.
– They are also called van der Waal
forces (although there is also a
repulsive van der Waal force)
– They are responsible for many physical
properties, including MPt and BPt
Intermolecular Forces
– There are several types DEPENDING
on the MOLECULAR POLARITY!
– Here are the 3 Main Types
• Ion-Dipole (Ch 12)
• Dipole-Dipole (and H-Bonding)
• London Dispersion
Dipole Dipole Forces
• Attractive force between POLAR molecules (have a dipole
moment)
• Electrostatic attraction of partial positive end of
molecule to partial negative end of another molecule
• In liquid or solid, molecules align themselves so are
attracted to several other molecules
• These dipole-dipole forces are much weaker than a real
covalent bond, about 3-4kJ/mol.
• So they may be broken with a low amount of energy, so
the solid melts, and the liquid evaporates!
431 kJ/mol
16 kJ/mol
Strength of Dipole-Dipole Forces
• There are 2 main factors in the strength of
dipole-dipole forces:
– Distance between molecules (the closer they
are, the stronger the dipole-dipole forces)
– Molecular Polarity (the more polar, the
stronger the dipole-dipole force)
Hydrogen Bonding: A special case of
dipole-dipole forces
• Some molecules have such a strong dipole-dipole
force, that this extra-strong dipole force was
given its own name: Hydrogen Bonding
• It is NOT really a bond, it is just an extra
strong dipole-dipole force.
• It occurs under certain circumstances.
H-Bonding
• Occurs when a molecule is small and very polar
• Occurs when have a N-H, O-H, or H-F bond
• The molecules have a very large dipole moment,
and they can get very close to one another due
to the small size of H, N, O, F
• H-Bonding is responsible for the very high
melting and boiling point of water
• H-Bonding is responsible for shape of proteins
and DNA
London Dispersion Forces
• But CO2 is a nonpolar molecule, so
how can it ever be a solid?
• There is another force, called the
London Dispersion Force.
• The London Force acts upon ALL
molecules, polar or not!
Loondon Dispersion Forces
• But London Forces are the ONLY
intermolecular force that operates
in nonpolar molecules.
• So how does it work when there are
no dipoles to create an attraction?
• Instantaneous, induced dipoles!
London Dispersion Forces
• London Forces seem weak, but they
can be very important, and many
nonpolar molecules are solids or
liquids.
• Waxes, oil, gasoline are examples.
London Dispersion Forces
• There are 2 factors in the strength
of London Force:
– Size and mass of molecule or atom
– Shape of molecule
Effect of Size/Mass on London
Dispersion Forces
• The larger the size or mass of a molecule, the
more electrons it has.
• These electrons are also generally further from
the nucleus.
• So the electrons can be distorted or pushed
and pulled from one end of the molecule more
readily, creating a larger temporary dipole.
Effect of Shape on London Dispersion Forces
• The closer molecules can align, the stronger the
London Force, and the higher the melting and
boiling points.
• Bulky molecules with “branching” can’t get as
close, and so have lower London Forces and will
have lower MPt and BPt than an “unbranched”
molecule of the same mass.
Intermolecular Force Summary
•
London Forces act upon all molecules, but are most
•
Dipole-dipole forces are found in polar molecules and
•
H-bonding occurs when very small and very
important for larger, heavier molecules
depend on the size and polarity of the molecules
electronegative N, O, or F atoms are present along
with N-H, O-H, or H-F bonds
Intermolecular Force Summary
• Although they are weak compared to covalent bonds,
these intermolecular forces greatly affect the boiling and
melting points of a compound
• The stronger the forces, the higher the melting and
boiling points
• The stronger the forces, the more likely a compound will
be a solid or liquid
• They also mean more deviations of a gas from the Ideal
Gas Law!