Intermolecular Forces - slider-dpchemistry-11

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Intermolecular Forces
11 DP Chemistry
London Dispersion Forces
The temporary separations of
charge that lead to the
London force attractions are
what attract one nonpolar
molecule to its neighbors. This
is due to the constant
movements of electrons
Fritz London
1900-1954
Dispersion forces increase with the
size of the molecules due to the
presence of more electrons. They
are also known as Van der Waal
forces.
London Forces in Hydrocarbons
Dipole-Dipole
Attractions
Attraction between
oppositely charged
regions of neighboring
covalent molecules.
The negative and
positive regions of
these molecules
represent regions of
higher and lower
electronegativity,
respectively.
The water dipole
The ammonia dipole
Hydrogen Bonding
Bonding between
hydrogen and more
electronegative
neighboring atoms
fluorine, oxygen and
nitrogen
Hydrogen bonding in Kevlar, a strong polymer used in bulletproof vests is represented by the dashed lines.
Hydrogen
Bonding
in Water
Hydrogen Bonding between
Ammonia and Water
Evidence for strength of H-bonding
HF and HCl are
similar molecules and
both have polar
bonds
However, their boiling points are very different
indicating much greater IM forces between HF.
Boiling points (0C)
HCl
HF
-85
20
Notice this trend is the
same for all hydrides
containing N, O, F.
Another Example
Any molecule which has a hydrogen atom
attached directly to an oxygen or a
nitrogen is capable of hydrogen bonding.
Such molecules will always have higher
boiling points than similarly sized
molecules which don't have an -O-H or
an -N-H group
Ethanol, CH3CH2-O-H, and
methoxymethane, CH3-O-CH3, both have
the same molecular formula, C2H6O.
Ethanol
Methoxymethane
78.5°C
-24.8°C
Referring to the structure,
explain why the bp difference
ANSWER:
Van der Waals forces are similar - they
have the same number of electrons and a
similar length.
However, ethanol has a hydrogen atom
attached directly to an oxygen with two
lone pairs as in a water molecule. Hydrogen
bonding can occur between ethanol
molecules, although not as effectively as in
water. The hydrogen bonding is limited by
the fact that there is only one hydrogen in
each ethanol molecule with sufficient +
charge.
In methoxymethane, the lone pairs on the
oxygen are still there, but the hydrogens
aren't sufficiently + for hydrogen bonds to
form. Except in some rather unusual cases,
the hydrogen atom has to be attached
directly to the very electronegative
element for hydrogen bonding to occur.
Relative magnitudes of forces
The types of bonding forces vary in their
strength as measured by average bond
energy.
Strongest
Covalent bonds (400 kcal)
Hydrogen bonding (12-16 kcal )
Dipole-dipole interactions (2-0.5 kcal)
Weakest
London forces (less than 1 kcal)
Exercises
1. H20 and H2S have the same molecular shape.
However, they have very different boiling
points (100 and -600C respectively) Draw
these molecules and explain the difference in
b.p. values.
2. What would you predict about the relative
b.p. values of NH3 and PH3?
3. Compare the relative strengths of IM forces
for the three compounds below (propane,
ethanal and ethanol).

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