Intermolecular Forces
Section 3.4
Pg. 105-117
1)
Explain intermolecular forces, London
(dispersion) forces, dipole-dipole attractions
and hydrogen bonding
2) Relate properties of substances to the predicted
intermolecular bonding in the substance.
BACKGROUND
• All chemical changes (reactions) are accompanied by energy
changes
▫ Energy is mostly heat, light, or electrical energy
▫ Energy can be released slowly (battery) or quickly (fireworks)
▫ Two types of energy changes are possible:
 EXOTHERMIC – energy is released into the surroundings
- the product’s bonds have less energy than the reactant’s bonds
 ENDOTHERMIC – energy is absorbed from the surroundings
- the product’s bonds have more energy than the reactant’s bonds
▫ Bond Energy – the energy required to break a chemical bond or the
energy released when a bond is formed
BACKGROUND
• There are three types of forces in matter:
1) Intranuclear force (bond) – bonds within the nucleus between protons and
neutrons (very strong)
2) Intramolecular force (bond) – bonds between atoms within the molecule or
between ions within the crystal lattice (quite strong)
3) Intermolecular force (bond) – bonds between molecules (quite weak); are
electrostatic (involve positive and negative charges)
There are 3 types of intermolecular bonds:
Weakest
a) London Force (a.k.a. London Dispersion Force, Dispersion Force)
Medium
b) Dipole-Dipole Forces (a.k.a. Polar Forces)
Strongest
c) Hydrogen Bonding
Note: “Van der Walls force” – includes London and dipole-dipole forces
1) London Force
• Simultaneous attraction between a momentary dipole in a
molecule and the momentary dipoles in surrounding molecules
momentary dipole: an uneven distribution of electrons around a
molecule, resulting in a temporary charge difference between its ends
They last for just
the instant that
the electrons are
not distributed
perfectly even.
1) London Force
• Fritz London also showed that momentary dipoles occurring in
adjacent molecules would result in an overall attraction
• The strength of the London force is directly related to the number
of electrons in the molecule, and inversely related to the distance
between the molecules.
▫ Increase electrons = Increase force (directly related)..
▫ Increase distance = Decrease force (inversely related)
1) London Force
• The key point is that:
▫ the more electrons a molecule has, the more easily
momentary dipoles will form, and the greater the
effect of the London force will be.
• London forces are present between all molecules,
whether any other type of attraction is present.
2) Dipole-Dipole Force
• The simultaneous attraction between oppositely
charged ends of polar molecules.
▫ Simply put, the attraction between diploes
Dipole: a partial separation of positive and negative charges
within a molecule, due to electronegativity differences
▫ Dipole-dipole forces are among the weakest
intermolecular forces, but still control important
properties (i.e. Solubility because water is polar))
2) Dipole-Dipole Force
In a liquid, polar molecules
can move and rotate to
maximize attractions and
minimize repulsions. The
net effect is greater overall
attraction.
The strength of the
dipole-dipole force is
dependent on the
overall polarity of the
molecule
Note: If a molecule is
polar it will be soluble
in water? Why?
2) Dipole-Dipole Forces
In a liquid:
In a solid:
Why do we care about intermolecular forces?
• We can use Dipole-Dipole and London Forces to predict Boiling Points
Compound (at SATP)
Electrons
Boiling Point (°C)
CH4(g)
10
-164
SiH4(g)
18
-112
GeH4(g)
36
-89
SnH4(g)
54
-52
A higher boiling point temperature means more energy has to be
added, thus we assume the intermolecular forces are stronger.
(see Learning Tip pg. 109)
Remember (if all other factors are equal):
1) The more polar the molecule = The stronger the dipole-dipole force
2) Increase the number of electrons = Increase the strength of London Force
Example #1
• Use Intermolecular force theory to predict which of the
following hydrocarbons has the highest boiling point:
▫ methane (CH4), ethane (C2H6), propane (C3H8), butane (C4H10)
1) Are the molecules polar or non-polar?
non-polar (no dipole-dipole force)
2) Which has more electrons?
butane: greatest # of e-’s = greatest London force
Check:
Alkane
Boiling Point (°C)
methane
-162
ethane
-89
propane
-42
butane
-0.5
Example #2
• Use Intermolecular force theory to predict which of the following has the
highest boiling point:
▫ bromine (Br2 )
1)
or
iodine monochloride (ICl)
Which has more electrons?
They are isoelectronic: have the same number of electrons (70 e-’s)
-Therefore the London force is the same (or nearly the same)
1)
Are the molecules polar or non-polar?
-Bromine is non-polar (has no dipole-dipole force; only London forces)
- Iodine monochloride is polar (has dipole-dipole forces and London forces)
- This extra attraction between ICl molecules produces a higher boiling point
Check:
Substance
Electrons
Boiling Point (°C)
bromine
70
59
iodine monochloride
70
97
You cannot predict boiling points when:
• One molecule has a stronger dipole-dipole force and the
other has a stronger London force
• The two molecules differ significantly in shape
• The central atom of either molecule has an incomplete octet
Practice
• Pg. 109 #1-4
• 4) a) Boron is stable with 6 valence eb) chloromethane (CH3Cl)
3) Hydrogen Bonding
• Occurs when a hydrogen atom bonded to a strongly
electronegative atom, (N, O and F) is attracted to a
lone pair of electrons in an adjacent molecule.
▫ Hydrogen nucleus (proton) is simultaneously attracted to
two pairs of electrons; one closer (in the same molecule)
and one further away (on the next molecule)
Why do you need a strongly
electronegative atom?
It pulls the hydrogen’s
electron away making it
“unshielded”, so the lone
pair on the other side can
come much closer
3) Hydrogen Bonding
• Hydrogen bonds are momentary attractive
forces between passing mobile molecules but are
the strongest of the intermolecular forces.
• Hydrogen bonds only act as continuous bonds
between molecules in solids, where the
molecules are moving slowly enough to be
locked into position.
• Hydrogen force would have been a better name.
3) Hydrogen Bonding
• In ice, hydrogen
bonds between the
molecules result in a
regular hexagonal
crystal structure.
• The ···H– represents
a hydrogen nucleus
(proton) being shared
unequally between
two pairs of electrons
3) Hydrogen Bonding
• Do lakes freeze from the
bottom-up or the top-down?
• Top–down, because water is
unique in that its solid form
(ice) is less dense than its
liquid form. Why??
• The hydrogen bonds hold
water molecules in a
hexagonal lattice with open
space in the center, which
explains the low density
(mass/volume) of ice.
Hydrogen Bonding in DNA
• FYI: The double helix of the DNA
molecule owes its unique structure
largely to hydrogen bonding.
• The red bonds are hydrogen bonds.
• If the helix were held together by
covalent bonds, the DNA molecule
would not be able to unravel and
replicate and life could not
continue!!
Why do we care about intermolecular forces?
• Explains surface tension, shape of a
meniscus, volatility and capillary action
1) Surface Tension
▫ Molecules within a liquid are
attracted by other molecules in all
directions equally, but right at the
surface, molecules are only attracted
downwards and sideways. This
means the net pull is downward so
the surface tends to stay intact
▫ The stronger the intermolecular force
the stronger the surface tension.
This shows water
adhering to the faucet
gaining mass until it is
stretched to a point
where the surface
tension can no longer
bind it to the faucet.
It then separates and
surface tension forms
the drop into a sphere.
Why do we care about intermolecular forces?
2) Capillary Action – due to adhesion (attraction between unlike
molecules) and cohesion (attraction of like molecules)
▫ The adhesion between water and glass is greater than the
cohesion between water molecules.
▫ The cohesion between mercury molecules is greater than the
adhesion between mercury and glass
Hg clip
Meniscus
In a sense, water
is pulled up the
tube by the
intermolecular
forces between
water and glass
Practice
• Pg. 117 # 1, 2, 4, 5
▫ #1 – use pg. 99 table to determine polarity
▫ #1 – look for NH2, NH, OH2, OH, to determine if hydrogen bonding is
possible
• Ex. CH3CHOHCH3 will it have hydrogen bonding?