Dipoles – Electronegativity

00:02 So let’s go back to covalent chemistry because it’s not just about what’s happening inside the molecule in terms of where the electrons are distributed but also how that distribution influences the macroscopic behaviour of those molecules.

00:16 So I talked to you earlier about the idea of oxygen or indeed other electronegative elements pulling electron density of a bond – pulling electrons away from the other atom by virtue of being more electronegative than the atom to which it’s attached.

00:34 And here we have an example of a very simple dipole. This is our carbon-oxygen, single-bonded system. Note I’ve drawn the arrow there which appears to be a positive charge with an arrow coming out of the end. And this is a way of trying to explain that, where we have that positive charge, we have something called delta positive. This is to say, because the electrons are mostly towards the oxygen end in this single bond, we create a partial positive charge on the carbon. By doing so, of course, that means we must create a partial negative charge on the oxygen. And this gives us a dipole.

01:15 And the dipole can be written, for example in the case of oxygen bound to hydrogen via a sigma molecular orbital as polar covalent. This is where we have an unequal distribution of charge along that bond. We have more electron density – or a greater probability of finding an electron nearer the more electronegative atom than we do from the least electronegative atom in that system.

01:42 And you can draw it one of two ways: either with the arrow shown at the bottom of the board with the positive charge or by using the term delta negative and delta positive, delta being very small in this case. Electronegativity can give reasons as to why chemical reactions occur but also can help to explain some of the macroscopic realities of some of those molecules by virtue of dipole-dipole interaction, as we’ll see later.

02:13 It is possible on a single molecule to have multiple dipoles. And a simple example here is water.

02:21 Shown here at the centre of the board we have the molecular structure for water: one oxygen, two hydrogens. Note the sigma molecular bonds between the oxygen and the hydrogen on both sides. We haven’t shown those two pairs of non-bonding electrons. But what you’re aware based on what you saw in the previous slide is that most of the electron density is concentrated around the oxygen at that part of the bond. And so what you get is a dipole. You get a distribution of charge, or a charge separation. So you have a delta negative on the oxygen and a delta positive on the hydrogen. And, if you think about it or if you try to imagine how this correlates to how water exists on this planet, it exists as a liquid. It doesn’t exist as a gas. And the reason it exists as a liquid is by virtue of the fact that you can get a delta positive interacting with a delta negative.

03:16 Opposites attract, likes repel. So the fact that you could have opposite partial charges forming electrostatic interaction with each other is one of the reasons why the temperature on the planet is not, normally, enough for you to form the gas of water. It stays as a liquid at low temperature. And you require more energy to break those bonds in that dipole-dipole interaction in order to overcome those that bring them together.

03:49 This concept can be widened to look at electron distribution in whole molecules. And there are two ways in which chemical groups can alter their distribution of electrons. One is inductive and one is known as resonance, or mesomeric, effects. Both of them are, to all intents and purposes, to do with the stability of the molecules that we are observing.

04:13 Inductive effects: these are similar to the electronegativity effect of single atoms as we showed. Functional groups, as a whole, can either be electron donating or they can be electron withdrawing. So, when we’re talking about atoms, it’s relatively easy to determine what they are. We look at where they are in the periodic table. If they’re fluorine, chlorine, iodine they tend to pull electron density away. They polarise the bond.

04:42 And they can be regarded as electron withdrawing. More electron-donating groups, you’d obviously have to look further to the left on the periodic table.

04:51 But it is in organic chemistry a little bit more complicated than that because you can actually have complex groups which serve either as electron-donating or electron-withdrawing species.