Introduction – Electronegativity

00:02 In the previous lecture, we talked about the affinity of nuclear positive charge for electrons.

00:10 We talked about it in the context of ionic bonding and we also talked about it in the context of covalent bonding. And this relates to the Pauling scale that we alluded to earlier as well.

00:23 Here is just a quick revision of that scale. If you look at the periodic table representation shown here on the board, you’ll see that the increase in electronegativity follows an increase in atomic number running from group I to Group VII and that increase in electronegativity is associated with the smallest number of shells running from the highest period to the lowest.

00:49 Large atoms have a low affinity because the electrons are a long way from a positive nucleus.

00:56 The core electrons shield the outer electrons from the nuclear charge. And, going across the periodic table, the affinity for bonding electrons increases because nuclear charge increases without a significant change in the distance of the electrons from the nucleus.

01:16 And this is okay in explaining how we observe bond polarisation in covalent bonds and also where we see ionic bonds forming.

01:26 Fluorine is the most electronegative atoms with a value of 4.0 on the Pauling scale.

01:34 And remember that Group 18, or Group 8 depending on your periodic table elements, do not react: they are the noble gases and therefore their shells are already full and they are unreactive.

01:47 All other elements have smaller values on the Pauling scale. And in bonds between atoms of different elements, the bonding electrons can be shared unequally.

01:58 Let’s have a look for example at carbon bonded to oxygen. Note I’ve shown you a sigma bonding: a single carbon-to-oxygen bond.

02:09 If we look at the individual values on the Pauling scale for carbon and for oxygen, we see they are 2.5 and 3.5 respectively. The difference between the two is 1. And what that means is we have a polar covalent bond. What that means is that the probability density or the chance of finding an electron within the sigma bond is greatest when we move towards the more electronegative element, in this case the oxygen. This is so called polarised.

02:45 It also – the scale – allows us to determine whether or not we’re likely to see an ionic or a covalent interaction.

02:52 If we look at one of the carbon-hydrogen sigma bonds of our friend methane from previous lectures, you’ll see that the difference in electronegativity between the two is only 0, based on 2.5 – 2.1. This gives us a relatively non-polar covalent bond where the electrons, or the chances of finding the electrons, are likely to be mostly in the centre between the two.

03:22 If on the other hand we look at the interaction between sodium and chlorine and we form an ionic sodium and chlorine bond, we see the difference is quite different: chlorine has an electronegativity on the Pauling scale of 3, sodium of 0.9. The difference between the two is 2.1 and the electrons, because they are much more strongly attracted to the chlorine, result in the formation of an ionic compound with formal loss of an electron from the sodium to the atomic orbital of the chlorine. So, in sodium chloride, as when you see other interactions between Group-1 and Group-2 and Group-6 and 7 atoms, you will observe an ionic compound being formed. Bear in mind ionic compounds exist as formula units. You would never ever refer to a salt or ionic compound as a molecule. Please be aware of that distinction.