So, let’s have a look at some of those intermolecular
Now, these should be, at least the first two,
should be those which you are already familiar
with. This is the idea of sharing electrons
in a covalent bond. As you can see, it has
the largest energy in terms of moles of bonds
at 300 to 450 kilojoules per mole. So, this
type of bonding is very, very strong.
Ionic bonding, we’ve come across before
as well, where we have a discrete unit charge
of a cation interacting with a discrete unit
charge of an anion. Because this is non-directional,
the approximate energy in kJ/mol is between
20 and 40.
The next important intermolecular interaction
is the ion dipole which gives us an approximate
energy of up 150 kJ/mol. Hydrogen bonding
was introduced in the latter part of Module
III and this is where we have this large dipole
occurring because of the electronegative effect
of oxygen pulling electrons in from the hydrogen-oxygen
sigma molecular orbital making the hydrogen
delta positive and the oxygen delta negative.
With this type of intermolecular interaction,
it’s possible to achieve approximate energy
of bonds of 37 kJ/mol.
Then we have weaker intermolecular forces:
dipole-dipole, such as, for example, as we’ll
see the interaction between two carbonyl groups,
hydrophobic interactions and van der Waals
So, let’s have a look at the weakest. Van
der Waals bonding exists between all atoms.
And this arises because the electron cloud
associated with an atom or molecule is constantly
moving. So, remember what we said originally
right back at the beginning of this entire
course in Module I. We talked about the idea
not of electrons flying around a nucleus in
a planetary fashion, but rather than being
found in a specific point or rather having
them there as a chance probability of finding
So, what I mean by that is that, when we’re
looking, for example, at a sigma orbital and
we show the electrons being in the centre
when you’ve got atoms which are of the same
type, what we actually mean is that there’s
a 95% probability of finding it there. But,
there’s also, obviously, by that definition,
a 5% possibility of finding an electron somewhere
else. And this is what’s happening here,
since we are looking at it from a time perspective,
we can never be certain at any given point
in time where precisely all of the electron
density happens to be. All we can say really
is the average.
So, what this means is that we may end up,
in the case of a small otherwise non-dipole-based
molecule, such as, for the sake of argument,
a bromine molecule, with a small instantaneous
dipole caused by the, if you like, almost-sudden
movement of charge from one end of the molecule
to the other at a given point in time.
This creates an artificial, very small dipole
moment where you’ve got more electron density
on one side of the molecule than you have
on the other. The net result being, you have
delta-negative and delta-positive dipoles.
Thus, these behave like small magnets and
can attract each other. And it’s van der
Waals’ forces that, for example, are responsible
for keeping together the atoms and the noble
gases together when they’re in gas form.
The only interactions they can exert against
each other is van der Waals.
Van der Waals forces though are very, very
weak. The larger the surface area and the
larger the number of electrons in the molecule,
the larger the interaction will be. And you
see this in the example of the melting and
boiling points of our halogens. If you run
down from fluorine to chlorine to bromine
to iodine, you’ll see that their actual
state at room temperature and pressure are
gas, gas, liquid, solid. As you go down, you’re
increasing the amount of charge that’s available
because, obviously, you’re moving down the
shells. Iodine has far more electrons in its
shells than fluorine does. And so, therefore,
the potential for a van der Waals interaction
with another molecule of iodine is so much
greater. And this results in them being together
more strongly and therefore, being a solid
at room temperature requiring more energy
to break them apart.
These interactions will only occur between
molecules which are very close together, around
0.4 to 0.6 nanomolars apart. The force, as
you can see here by this equation, drops off
quickly as the molecules move apart. And,
indeed, the force is proportional to 1/d to
the 6, where d is the distance between them.
These forces are insignificant for individual
atoms, but can be important in terms of pairs
of molecules with lots of atoms, especially
if the surfaces of the molecules are the right
shape to allow a close fit.