Types of Interaction – Biological Interactions

00:00 So, let’s have a look at some of those intermolecular interactions.

00:05 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.

00:27 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.

00:43 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.

01:22 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 interactions.

01:38 So, let’s have a look at the weakest. Van der Waals bonding exists between all atoms.

01:45 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 them there.

02:08 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.

02:41 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.

03:05 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.

03:32 The only interactions they can exert against each other is van der Waals.

03:36 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.

04:31 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.

04:52 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.