INTER-MOLECULAR
FORCES
State of matter?
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)
have high melting and boiling points
 imf not applicable because NO
MOLECULES exist
SIMPLE molecules (eg H2O, H2, CH4 etc)
have much lower melting and boiling points
 imf 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 provide good
indications of variations in IMF
Don’t forget !!!!
Strong 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.
Van der Waal forces
Dipole-dipole forces
Hydrogen bonds
For similar size
molecules, imf
strength INCREASES
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
H 2Se
246
H 2Te
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
VW Forces only
Hydrogen bonds
DP-DP +VW forces
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 mol.
δ+ H-F:
δ+ H-F:
δ+ H-F:
H2O has approx. two H bond per mol.
..
δ+ H-O:
H
δ+
..
δ+ H-O:
H
δ+
..
δ+ H-O:
H
δ+
Hence water’s unusually HIGH mpt and bpt
NH3 has approx. one H bond per mol.
H
δ+ H-N:
H
H
δ+ H-N:
H
H
δ+ H-N:
H
In LIQUID,
H-bonds are
continuously
breaking and
reforming
In SOLID,
H-bonds are
permanent
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
H 2S
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’s molecules are DIFFICULT (require
a lot of added energy) to separate
Hence water’s unusually HIGH boiling point (100ºC)
And melting point (0ºC)
when compared to other molecules
of similar size / mass
eg H2S (a heavier molecule!) is a GAS at
room temperature
Also because of it’s strong H-bonding, water
has an unusually H IGH SURFACE TENSION
This creates a “skin-like” effect on the surface of water
This allows insects like the water
strider to “walk on water”!
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 during freezing of
water can burst pipes
and ice floats on water
Dipole-Dipole Forces
Dipoles in molecules arise from uneven
electron distributions
caused by differences in the
electronegativity of the atoms in the
molecule.
e.g.1 Hydrogen chloride:
+
H
-
Cl
+
H
-
Cl
H < Cl in electronegativity
 bond e- drawn towards Cl
 permanent dipole in molecule ;
δ+ one end, δ- at other
 dipole-dipole force between δ+ of one
mol and δ- of neighbouring mol
e.g.2 Hydrogen sulphide:
CH
+
CH
CH
-
S
O
+
-
CH
H < S in electronegativity
 permanent dipole
 dipole-dipole force
S
O
NB WEAKER van der Waal
forces ALSO occur
whenever dipole-dipole
forces occur
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-
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-
e-
e-
e-
e-
-
+
eee-
e-
ee-
e-
-
Van de Waals forces get stronger as
number of electrons in molecule increases
 Boiling points of noble gases INCREASE down group
as e- shells are added
He
Ne
Ar
Kr
2e-
10e-
18e-
36e-
Increasing boiling point
Similarly, bpt group IV hydrides (CH4, SiH4 etc)
INCREASE down group as Mr increases
Consider the data for the following non-polar molecules:
Pentane
C5H12
Bpt = 36oC
2,2 Dimethylpropane
Also C5H12
Bpt = 10oC
SAME Mr but much stronger VW forces for PENTANE!
Straight chain molecule
CH3-CH2-CH2-CH2- CH3
Branched chain molecule
CH3
CH3-C- CH3
 mols can get close
CH3
 mols get LESS close
 stronger VW forces
 weaker VW forces
 higher bpt
 lower bpt
The End