Intermolecular Forces
Topic 4.3
Intermolecular Forces
• Intramolecular forces – refer to the forces that hold
atoms together within molecules or formula units; a
chemical bond.
• Intermolecular forces- refer to the forces between
particles in a substance. These forces are very weak
compared to intramolecular forces.
• For example, 927 kJ of energy is required to
decompose one mole of water vapor into H and O
atoms. However, only 40.7 kJ are required to convert
one mole of liquid water into steam at 100 oC.
Intermolecular Forces
Intermolecular Forces
• If it were not for the existence of intermolecular
attractions, condensed phases (liquids and solids)
would not exist. These are the forces that hold the
particles close to one another in liquids and solids.
• Changing the state of a substance from solid to liquid
or liquid to gas involves separating particles by
overcoming the forces between them.
• The stronger the intermolecular forces, the more
energy needed to separate the particles in substance,
and so the higher the substance’s melting and boiling
points
Intermolecular Forces
Intermolecular Forces
• We will consider three types of intermolecular
forces:
– London Dispersion forces
– Dipole-Dipole attraction
– Hydrogen bonding
London Dispersion Forces
• Non-polar molecules such as chlorine (Cl2) have
no separation of charges within their bonds
because the electrons are shared equally.
• However, because electron density at any one
moment may be greater over one atom than the
other, some separation of charge may occur, this
is known as a temporary or instantaneous dipole.
• This instantaneous dipole may influence the
electron distribution in the neighboring molecule
by causing an induced dipole.
London Dispersion Forces
• As a result, weak forces of attraction, known
as London dispersion forces, will occur
between opposite ends of these two
temporary dipoles in the molecules.
• London dispersion forces are present between
all types of molecules, but are the only kind of
intermolecular forces present among
symmetrical non-polar substances and
monatomic species such as noble gases.
London Dispersion Forces
• They are the weakest form of intermolecular
forces.
• Their strength increases as the number of
electrons within a molecule increases, that is,
with increasing molecular mass.
• Substances that are held together by van der
Waals’ forces generally have low melting and
boiling points, because relatively little energy is
required to break the forces and separate the
molecules from each other.
London Dispersion Forces
• Boiling point data show how the strength of
van der Waals’ forces increases with
increasing molecular mass
Element
Mr
Boiling Point/oC
State at room
temperature
F2
38
-188
Gas
Cl2
71
-34
Gas
Br2
160
59
Liquid
I2
254
185
Solid
Dipole-Dipole Attraction
• Molecules such as HCl have a permanent
separation of charge within their bonds as a
result of the difference in electronegativity of the
atoms. The Cl end has a partial negative charge
(δ-) while the H end has a partial positive
charge(δ+).
• This is known as a permanent dipole, and results
in opposite charges in neighboring molecules
attracting each other, generating a force known
as a dipole-dipole attraction.
Dipole-Dipole Attraction
Dipole-Dipole Attraction
• The strength of the dipole-dipole attraction
depends on the degree of polarity within the
bond. It will decrease as the degree of polarity
with the molecule decreases.
• For example, it will decrease in strength in the
order: HCl>HBr>HI.
• Dipole-dipole forces are stronger than London
dispersion forces, consequently, the melting and
boiling points of polar compounds are higher
than those of non-polar substances of
comparable molecular mass.
Hydrogen Bonding
• When a molecule contains hydrogen
covalently bonded to a very electronegative
atom (fluorine, nitrogen, or oxygen), these
molecules are attracted to each other by a
strong type of intermolecular force called
hydrogen bonding.
• The hydrogen bond is in essence a particular
case of dipole-dipole attraction.
Hydrogen Bonding
Hydrogen Bonding
• The large electronegativity difference between
hydrogen and the bonded fluorine, oxygen, or
nitrogen, results in the electron pair being
pulled away from the hydrogen. Given its
small size and the fact that it has no other
electrons to shield the nucleus, the hydrogen
now exerts a strong attractive force on a lone
pair of electrons in the electronegative atom
of a neighboring molecule. This is the
hydrogen bond.
Hydrogen Bonding
Hydrogen Bonding
• Hydrogen bonds are the strongest form of
intermolecular attraction.
• Consequently, they cause boiling points of
substances that contain them to be
significantly higher than would be predicted
from their molecular weight.
Hydrogen Bonding
• In all four groups there is an observable trend of boiling point increasing down
the group as molecular weight increases. The anomalies are NH3, HF and H2O
which have much higher boiling points than expected from their molecular
weight.
•T his can only be explained by the presence of hydrogen bonding in these
molecules
Hydrogen Bonding
• Likewise, when we compare the boiling points
of some organic molecules that similar or
equal molecular weights, we find a higher
value where hydrogen bonding occurs
between the molecules.
• For example:
CH3-O-CH3
CH3-CH2-O-H
MW= 46
MW= 46
Does not form hydrogen bonding
Boiling point = -23oC
Forms hydrogen bonds
Boiling point = 79oC
Relative Strength of
Intermolecular Forces
• H-bonding > dipole-dipole > London
• With molecules possessing similar molecular
weights, the molecule having the stronger IMF
will usually have the larger bp/mp. This is
especially true when the IMF is H-bonding.
Relative Strength of
Intermolecular Forces
• With molecules in which one of the atoms has been
changed with another atom within the same group,
the molecule having the larger formula weight will
usually have the larger bp/mp, except when the IMF
is H-bonding.
• Thus, NH3 (IMF – H-bonding) will have a higher
bp/mp than PH3 (IMF- dipole-dipole) even though
NH3 has a smaller molecular weight.
• However, AsH3 will have a higher bp/mp than PH3
because both have the same IMF, but AsH3 has a
larger molecular weight.
Hydrogen Bonding
• Water can form up to four hydrogen bonds
because it has two lone pairs on the oxygen
atom. Liquid water will contain fewer than this
number, but in the solid form, ice, each water
molecule forms 4 hydrogen bonds.
• The result is a tetrahedral arrangement that holds
the molecules a fixed distance apart, forming a
fairly open structure which is less dense than
liquid water, allowing ice to float in water.
Hydrogen Bonding