In the previous lecture, we talked about the
affinity of nuclear positive charge for electrons.
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
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
Large atoms have a low affinity because the
electrons are a long way from a positive nucleus.
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.
And this is okay in explaining how we observe
bond polarisation in covalent bonds and also
where we see ionic bonds forming.
Fluorine is the most electronegative atoms
with a value of 4.0 on the Pauling scale.
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.
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.
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.
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.
It also – the scale – allows us to determine
whether or not we’re likely to see an ionic
or a covalent interaction.
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
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.