Title: Lesson 7 Intermolecular Forces
Learning Objectives:
– Learn to identify and explain the three types of intermolecular forces:
• Van der Waals
• Permanent dipole-dipole
• Hydrogen bonds
– Understand and explain the effects of the above on melting/boiling points
Refresh

Use the VSEPR theory to deduce the shape of H3O+ and
C2H4. For each species, draw the Lewis structure, name
the shape, and state the value of the bond angle(s).
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Note Taking


Do not copy the notes, re-express them
Include diagrams
Van der Waals / Temporary
Dipole-Induced Dipole
Dipole-Dipole / Permanent
Dipole
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Hydrogen Bonds
Intermolecular Forces

The attractive forces between
molecules

It is these that are partially broken
during melting, and fully broken
during boiling

Note: when molecular compounds
melt/boil, the bonds in the
molecule do not break, it is just the
attractive forces between the
molecules that break
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Intermolecular Forces (imf)
These are weak electrostatic forces of
attraction between neighbouring molecules.
They are much weaker than covalent, ionic or
metallic bonding.
They influence ONLY the physical properties
of molecules.
GIANT structures
covalent (eg diamond), or ionic (eg NaCl) or
metallic (eg Cu)
have high melting and boiling points
 imf not applicable because NO
separate MOLECULES exist
SIMPLE molecules (eg H2O, H2, CH4 etc)
have much lower melting and boiling points
 imf are applicable
IMF influence PHYSICAL properties :

Melting points and boiling points

Solubility in water and other solvents

3D shapes of complex molecules such as DNA

Viscosity of liquids.

Density
etc etc
Boiling point variations are very good
indicators of variations in IMF
Don’t forget !!!!
Strong covalent bonds within molecules
are not broken when molecular
substances are vaporized
Weak imf between molecules
are broken when molecular
substances are vaporized
Boiling point INCREASE
IMF strength INCREASE
Boiling point variations suggest 3 types of imf :
1.
2.
3.
London (Dispersion)
Dipole-dipole forces
Hydrogen bonds
For similar size
molecules, imf
strength INCREASES
NOTE: Van der Waals’ forces is an umbrella term to cover
both London dispersion and dipole-dipole attractions...
GROUP
FORMULA FORMULA FORMULA FORMULA
& BPt /K
& BPt /K
& BPt /K
& BPt /K
IV
CH4
109
SiH4
161
GeH4
190
SnH4
221
V
NH3
240
PH3
185
AsH3
218
SbH3
256
VI
H2O
373
H2S
212
H2Se
246
H2Te
280
VII
HF
293
HCl
188
HBr
206
HI
238
PERIOD
2
3
4
5
Noble
Gases
He
20
Ne
Ne
27
27
Ar
Ar
87
87
Kr
121
London Forces only
Hydrogen bonds
DP-DP +L forces
London (dispersion) aka
Temporary or instantaneous induced dipole forces

Non-polar molecules such as Cl2, have no
permanent separation of charge (no permanent
dipole)...

However, random electron movements create a
small, temporary dipole

This induces a similar dipole in a neighbouring
molecule

This creates a small attraction between them

These are weak and exist only for the tiniest
fraction of a second

London (dispersion) forces are present in all
molecules


Increase with molecular mass
Decrease with the roundness of a molecule
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Van der Waal Forces
Consider non-polar molecules such as Ne, I2, CH4 etc
Electron cloud of the molecule is in constant random motion
Leading to momentary electron density imbalance.
Leading to a temporary dipole in the molecule
which induces a temporary dipole in neighbouring molecule.
Leading to momentary attraction between temporary
dipoles which IS the van der Waal force
e-
ee-
+
eee-
e-
ee-
e-
e-
-
+
eee-
e-
ee-
e-
-
Van der Waals forces
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The strength of
van der Waals
forces increases
as molecular size
increases.
This is illustrated
by the boiling
points of group 7
elements.
boiling point (°C)
Strength of van der Waals forces
200
150
100
50
0
-50
-100
-150
-200
F2
Cl2
Br2
element
I2
Atomic radius increases down the group, so the outer
electrons become further from the nucleus. They are
attracted less strongly by the nucleus and so temporary
dipoles are easier to induce.
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Strength of van der Waals forces
The points of contact between molecules also affects the
strength of van der Waals forces.
butane (C4H10)
2-methylpropane (C4H10)
boiling point = 272 K
boiling point = 261 K
Straight chain alkanes can pack closer together than
branched alkanes, creating more points of contact between
molecules. This results in stronger van der Waals forces.
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Boiling points increase as number of
electrons increases...
• More electrons mean a larger electron cloud density and
this will induce a stronger attraction between molecules...
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Boiling points of alkanes
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Dipole-dipole forces
aka
Permanent dipole forces


Different atoms have different electronegativities, which
means there will be variations in the electron charge
density in different parts of a molecule
-
If a molecule is not symmetrical, the variation produces
a dipole where a molecule has a positive and a negative
end


+
The end with high charge density is The end with low charge density is +

Oppositely charged dipoles attract each other.

This is a relatively strong attractive force

If a molecule is symmetrical, variations in electron
charge density cancel each other out and the molecule
is non-polar... (Think of the tug of war example!)
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-
+
-
Permanent dipole–dipole forces
If molecules contain bonds with a permanent dipole, the
molecules may align so there is electrostatic attraction
between the opposite charges on neighbouring molecules.
Permanent
dipole–dipole
forces (dotted
lines) occur in
hydrogen chloride
(HCl) gas.
The permanent dipole–dipole forces are approximately
one hundredth the strength of a covalent bond.
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Permanent dipole–dipole or not?
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Melting and Boiling Points are stronger
in Dipole-Dipole



Strength can vary depending on the distance and relative orientation of the
dipoles.
Generally stronger than London forces so energy needed to separate
bonds will be greater.
NOTE: Weaker London forces also occur alongside Dipole-Dipole
forces...
NOTE! It’s important to compare substances with a similar molecular mass –
otherwise the difference can be attributed to stronger London forces based on
more electrons...
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Van der Waals’ forces is an umbrella term
for:

I.e. van der Waals’ forces refers to all forces between
molecules that do not involve electrostatic attractions
between ions or bond formation.
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Why do molecules such as CCl4, BF3 and
BeCl2 NOT show dipole-dipole forces?
Individual bonds are polar eg δ+Be-Clδbut the molecules are NOT because
they are SYMMETRICAL
 bond dipoles CANCEL
 NON-POLAR molecule
δ-
δ+
δ-
Cl–Be–Cl
δ+
Fδ-
δ-F
B3δ+
Fδ-
Hydrogen bonds
H-bonds
aka

The strongest type of intermolecular force

They occur between a nitrogen, oxygen or
fluorine and a hydrogen that is bonded to a
nitrogen, oxygen or fluorine

N, O and F are the three most electronegative
elements, and all have lone-pairs when bonded

When H is bonded to N, O or F, the electrons
in the bonded are strongly attracted to the
N/O/F, leaving the H very positive

The lone pair on the N/O/F is strongly
attracted to the positive hydrogen
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What is hydrogen bonding?
When hydrogen bonds to nitrogen, oxygen or fluorine, a
larger dipole occurs than in other polar bonds.
This is because these atoms are
highly electronegative due to their
high nuclear charge and small size.
When these atoms bond to hydrogen,
electrons are withdrawn from the H
atom, making it slightly positive.
The H atom is very small so the positive charge is more
concentrated, making it easier to link with other molecules.
Hydrogen bonds are therefore particularly strong examples
of permanent dipole–dipole forces.
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Hydrogen bonding
In molecules with OH
or NH groups, a lone
pair of electrons on
nitrogen or oxygen is
attracted to the slight
positive charge on the
hydrogen on a
neighbouring molecule.
hydrogen
bond
lone pair
Hydrogen bonding makes the melting and boiling points of
water higher than might be expected. It also means that
alcohols have much higher boiling points than alkanes of a
similar size.
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Hydrogen Bonding
For these to occur you need:
1. A VERY electronegative atom with
an available lone pair of electrons,
ONLY F, O & N are sufficiently electronegative
and 2. a H atom directly bonded to a VERY electronegative atom
- electronegative atom draws e- away from H (de-shields it)
making it SLIGHTLY positive, δ+H
ONLY δ+H-F , δ+H-O or δ+H-N are appropriate.
A hydrogen bond = the attraction between a lone
pair on a N, O or F atom and a de-shielded H
atom in a δ+H-F, δ+H-N or δ+H-O bond
HF has approx. one H bond per molecule
δ+ H-F:
δ+ H-F:
δ+ H-F:
H2O has approx. two H bond per molecule
..
δ+ H-O:
H
δ+
..
δ+ H-O:
H
δ+
..
δ+ H-O:
H
δ+
In LIQUID,
H-bonds are
continuously
breaking and
reforming
In SOLID,
H-bonds are
permanent
Hence water’s unusually HIGH mpt and bpt
NH3 has approx. one H bond per molecule
H
δ+ H-N:
H
H
δ+ H-N:
H
H
δ+ H-N:
H
In GAS,
H-bonds are
completely
broken
δ+H-F: , δ+H-O:
, δ+H-N:
Decreasing strength of individual H bonds
because electronegativity decreases
but water forms TWO H-bonds per molecule
 order of b pt is H2O >> HF > NH3
not
HF > H2O > NH3
Further examples :
CH3CH2OH
will H-bond.
-O-Hδ+ - - - :O-
CH3-C-CH3
O
will not H-bond.
O bonded to C, not H
H2S
will not H-bond
H bonded to S which is
NOT electronegative
enough for H-bonds
Hydrogen Bonding and the Unusual Physical Properties of Water
H2O has approx. two H-bonds per molecule
..
δ+ H-O:
H
δ+
..
δ+ H-O:
H
δ+
..
δ+ H-O:
H
δ+
 for such SMALL molecules, water molecules are DIFFICULT
(require a lot of added energy) to separate
Hence water’s unusually HIGH melting point (0ºC)
and boiling point (100ºC)
when compared to other molecules
of similar size / mass
eg H2S (a heavier molecule!) is a GAS at room
temperature because it does not hydrogen bond
400
B Pt’s of NH3, H2O and
HF are UNUSUALLY high
 imf UNUSUALLY STRONG
 HYDROGEN BONDS
BPt (/K)
350
300
250
200
150
100
50
Noble gases
0
2
3
4
Period
5
Hydrogen bonding and boiling points
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Boiling points of the hydrogen halides
40
20
0
-20
-40
-60
-80
-100
HF
boiling point (°C)
The boiling point
of hydrogen
fluoride is much
higher than that of
other hydrogen
halides, due to
fluorine’s high
electronegativity.
HCl
HBr
The means that hydrogen bonding between molecules of
hydrogen fluoride is much stronger than the permanent
dipole–dipole forces between molecules of other
hydrogen halides. More energy is therefore required to
separate the molecules of hydrogen fluoride.
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HI
Permanent dipole–dipole forces
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Comparing Boiling Points...

If we compare these different forms of C2H6O (known as
isomers) we can see how the presence of H-bonds affects
the boiling point...
O is bonded to C so no H-bond
O is bonded to H so H-bond present
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H2O the anomaly...





Water has 2 hydrogen atoms and 2 pairs of lone pairs on the oxygen
Hence, it can form 4 hydrogen bonds with neighbouring water molecules.
Liquid water has fewer bonds, but ice uses up 4 bonds which results in a
tetrahedral shape that is fixed and open...
So ice is less dense than water as it expands...
Usually solids form closely packed particles and become more dense...
4 H-Bonds circled
Structure of ice with 4 H-Bonds
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Also because of it’s strong H-bonding, ice has an
unusually LOW DENSITY compared to water
In ICE, the maximum number
of H-bonds are operative
 molecules are held apart in a tetrahedral
arrangement of covalent and H-bonds
 much empty space between molecules
 larger volume than the same mass of water
 LOWER DENSITY than water
 expansion of water during
freezing can burst pipes
and ice floats on water
Also because of it’s strong H-bonding, water
has an unusually H IGH SURFACE TENSION
Hydrogen bonds “pull” molecules at the surface inwards
creating a “skin-like” effect on the surface of water
This allows insects like the water
strider to “walk on water”!
Summary of IMF
Covalent substance have a lower M.P and B.P compared to ionic because the
energy needed to overcome the IMF is lower than the energy needed to
break electrostatic attractions in an ionic lattice...
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Solubility
TIP! Think like for like...

Non polar substances can dissolve in non polar solvents by formation
of London dispersion forces between solute and solvent. E.g. non
polar Br2 can dissolve in non polar paraffin oil (a hydrocarbon)

Polar substances can dissolve in polar solvents e.g. water. Dipole
interactions and H-Bonding are responsible. E.g. HCl, glucose
(C6H12O6) and ethanol (C2H5OH) are polar substances that can
dissolve
Note: Larger
molecules where only
a small part is polar
will be less soluble as
the non polar parts
will not disassociate in
water.
Polar substances have a low solubility in non polar because
the Dipole-Dipole forces keep them together and not
interacting with the solvent...
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Electrical Conductivity



Covalent compounds do not conduct electricity (no ions)
An exception is HCl dissolved in water – it’s ions H+ and Cl- disassociate
in water.
Giant covalent molecules e.g. Graphite and Graphene are conductors
(mobile electrons). Fullerene and Silicon are semi conductors.
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Effects of intermolecular forces

Intermolecular forces play an important role in the
properties of compounds including:

Melting/boiling point:


Volatility:


Stronger intermolecular forces  higher m.p./b.p.
Stronger intermolecular forces  lower volatility
Solubility: like dissolves in like


Polar solutes dissolve best in polar solvents
Non-polar solutes dissolve best in non-polar solvents
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Summary of Physical Properties
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Looking into intermolecular forces

Complete the activity here to research and model
intermolecular forces
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Summary

Three types of intermolecular force, from strongest to
weakest:

Hydrogen bonds


Dipole-dipole


Between N/O/F and H attached to N/O/F
Between permanent dipoles on asymmetric molecules
London (Dispersion)

Between instantaneous dipoles formed on any molecule/atom
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