Hund's rule states that when you've got…
filling a degenerate orbital such as the 2p,
where we have three orbitals which have a
similar shape but just a different orientation
along the axis, that they would preferably
fill unpaired. This is the lowest energy configuration.
So if we take, for example… Let's take,
for example, this species here: carbon, okay?
There'll be two electrons... Sorry, boron.
Let's take, for example, boron, in this case,
where we have five possible electrons. As
indicated in the previous slide, we have two
electrons which fill the first principle quantum
shell (1s) and then we have two electrons
which will fill the next orbital in the second
shell. And we finally have one electron which
occupies one of the orbitals in the 2p subshell.
If we were to consider carbon, Hund's rule
would mean that the next electron (because
carbon possesses six electrons) would go unpaired
into the adjoining 2p orbital, because they
are of the same energy. It's only when all
orbitals of the same energy are singly filled
do you start to observe pairing of these orbitals.
And this is very important, because it is
unpaired electrons within individual atoms
which are responsible, to a large extent,
for the activity, or the reactivity, that
So for carbon, as I indicated, as we increase
the number of electrons, they go into the
unoccupied 2p orbitals. If we look, for example,
at oxygen, which is further along the periodic
table, it's only after we've complete… we've
actually added one electron to each individual
2p orbital that we start to pair up the electrons.
Individually occupied orbitals are, as you
can see here, of lower energy than paired
electrons in orbitals.
So let's have a look and break it down. As
you can see here, we're looking at principal
quantum number, which is shown at the beginning
of the tree, rising in energy: 1, 2, and 3.
Okay? If we break down those orbitals which
are available, we can consider those as subshells
of the principal quantum number. So in the
case of principal quantum number 1, we have
a single possible orbital: the s orbital.
And therefore, we have the designation 1s.
Maximum allowable electron configuration in
that case is 1s2. Then, we have the second
shell. Remember, we've unlocked the 2p orbitals.
Now, the maximum number of electrons
there is theoretically eight, because we can
have two electrons in the 2s subshell, and
we can have six electrons in the 2p subshell.
This gives us a total number of electrons
in the second shell of eight. Then we look
at the third shell, bearing in mind that we've
unlocked another set of orbitals, the diffuse
orbitals. So now it's possible to put two
electrons in the 3s, six electrons in the
3p, and 10 electrons in the 3d, as we'll come
to. Now, unfortunately, nothing's ever simple,
and as you can see here, we have the so-called
4s, which rather counterintuitively appears
to be of lower energy than the 3d. And this
ostensibly flies in the face of all we've
discussed about quantum mechanics. But the
reality is this: that the nature of the s
orbital in this current situation actually
means it's of lower energy initially than
the 3d, and it is actually populated with
electrons first. However, if you have a full
3d outer shell, the energy of the 4s is actually
forced up, and now, the principal quantum
number 4 gives a lowest energy s orbital,
which exists at higher energy than the 3d
orbital. And this chart shows the order of
filling that you should be familiar with.
Once you get to the transition metals, things
change slightly, and we'll talk about that
a little less… a little more later.
So now, based on your understanding, you should
be able to write out the electron configuration
of the first 20 elements in the periodic stable…
in the periodic table. Even if you can't necessarily
remember all of the chemical symbols, you
should be able to do this. And by understanding
what electrons are where, it's possible for
you to understand how they will react. The
configurations, as we've shown, refer to the
ground states of the atom for the element
in its atomic form. Bearing in mind, of course,
in ionic forms, as we'll see, if you lose
or gain electrons, the electron configuration
will also, therefore, change. Putting the
correct amount of energy in can lead to electrons
jumping to high energy levels and giving rise
to excited states, which we may touch upon
later. If we look here on the periodic table
extract, which I've shown you, which you can
see on this board, you can observe that I've
shown you hydrogen, lithium, beryllium, sodium,
and magnesium again. Hydrogen, as we indicated,
has electrons only in the 1s. Beryllium: The
valence electrons are in the 2s. And sodium,
magnesium: in the 3s.