Right. Now, let’s have a quick look at dipole-dipole
interactions. If you look on the right hand
side of the screen on the board, you’ll
see here two carbonyl compounds. Indeed, we
will call them ketones, in this particular
case. As you recall, there is a slight positive
charge conveyed by the electronegativity of
the oxygen in the carbonyl-carbon position.
The electron density being pulled by the oxygen
means there is a partial negative charge on
And what this results in is an interaction
between the partial negative charge on the
oxygen with the partial positive charge on
a neighbouring ketone. And this is an example
of a dipole-dipole interaction. And in many
molecules, you’ll find that there are permanent
dipoles due to the difference in electronegativities
of atoms sharing chemical bonds.
Let’s have a look again at carbonyl groups.
Note the delta positive on the carbon and
the delta negative on the oxygen. This is
the same not just for the carbonyl series,
but bear in mind, also for things like amides,
which contain carbonyl groups as well and
These dipoles, as I’ve said, will attract
one another leading to bonds which are stronger
than van der Waals. And this is an example
of an electrostatic interaction.
Right. Let’s move on now to hydrogen bonds.
You should already be familiar with the concept
of hydrogen bonds from Module III. And these
are special types of dipole-dipole interaction.
And this occurs when you’ve got hydrogen
directly bonded via a covalent bond to nitrogen,
oxygen or sulphur. So, those are the three:
nitrogen, oxygen or sulphur. And we haven’t
talked much about sulphur and we’re not
going to. We’re going to focus on nitrogen,
in the case of amines, and oxygen, in the
case of alcohols, and carboxylic acids.
As you can see from the diagram shown on the
right hand side, we show an idealised system
where the delta positive charge on the hydrogen
from that dipole interacts strongly with the
delta negative charge on the oxygen of a neighbouring
water molecule. And here you have the argument
of donor and receptor.
We consider something as an acceptor, if it
is the electron-rich component in an interaction
and we consider something as a donor, if it
is the electron-weak part i.e. that which
is accepting the electrons.
Two molecules of water are shown here and,
due the difference in electronegativity between
oxygen and hydrogen, the OH bond is, as we’ve
said, polarised. The lone pairs of electrons
on the oxygen bond with the hydrogen which
has a partial positive charge.
Okay. The hydrogen is effectively shared between
a donor and an acceptor in a hydrogen bond.
Which atom is it more tightly linked to though?
Donors can form one hydrogen bond for each
hydrogen. Acceptors, however, can form one
hydrogen bond for each electron lone pair.
Hydrogen bonds are, therefore, directional.
And so, the importance is not necessarily
just the distance per se, but also the angle
of attack. If you look at the bottom here,
you’ll see there are two diagrams. One actually
shows a hydrogen bonding of, in this case,
a secondary amine with a ketone. As you can
see, when the bonds are in line, when there
is no need to change direction and the angle
is 180 degrees over the top, you can see that
that’s a reasonably strong hydrogen-bond
interaction. However, when the orientation
is not directly in line, you have a weak bond.
So, these are the two examples of a strong
and weak hydrogen bond.
Hydrogen bonds are important not only in terms
of drug-target interactions, but as we’ve
seen before, holding together the structure
of both proteins and also DNA.
Now, let’s quickly have a look at ion-dipole
bonds. Many drugs, sorry, many drug molecules