00:02
So these energy levels as postulated by Niels
Bohr are called shells. And electrons, rather
than existing just as a random cloud at a
specific distance, exist at a number of different
distances from the central nucleus. The highest
energies are actually further away. The principle
quantum number is given as the actual lowest
possible quantum value for that shell and
is regarded as the principal quantum number.
So, if we look at the shell which is nearest
to the nucleus, we see it has the quantum
number 1. If we move one shell away, it has
the quantum number 2 then 3 and then 4. As
the shells become larger – i.e. as they
move further and further away from the nucleus
– we can see that they hold more and more
electrons. And we’ll be going through the
makeup of these shells in the next lecture
when we start looking at orbitals. But, from
a quantum perspective, each shell can carry
a specific number of electrons. Shell 1, which
is that which is nearest the nucleus, can
carry only 2 electrons. Shell 2 can carry
8, shell 3 can carry 18 and shell 4 can carry
32. And as we’ll see, this is because as
we increase the shell number – the principal
quantum number – we have at our availability
a large number of orbitals. And, as we’ll
see, each orbital can contain two electrons.
The relative energy of these shells increases
as we move further and further away. But as
I indicated in the previous slide, this means
that the relative energy required to remove
an electron from a shell actually decreases.
02:00
In order to move an electron to a high-energy
shell, which is analogous to the idea of a
person moving from one step to the other,
a specific amount of energy is required. And
that energy is called the quantum energy.
This is the energy – a small amount of energy
– which is required to move a particle from
a specific distance away from a proton to
the next distance away from the proton – i.e.
to move it from shell 1, for example, to shell
2. And as you’ll see when we start talking
about molecular orbitals interactions, these
movements of electrons from higher shells
to lower shells are the origins of things
like fluorescence and phosphorescence. Indeed,
if you look at the atomic absorption or atomic
emission spectroscopy for individual atoms,
you’ll see that their colours are to do
with the transitions of electrons from one
shell to the other and then back again.
02:57
So, in the case of our idealised atom which
is shown on the board here, you’ll see that
we have the movement of an electron from a
shell nearest the nucleus to a shell further
from the nucleus when a quantum of energy
is absorbed. Now this energy can be light
energy, for example, or it could theoretically
be heat energy.
03:22
When an electron moves back from its outer
shell, as shown here shown with the green line,
back to that electron which is nearest the
nucleus, we see a quantum of energy emitted,
usually as a photon of light.
So the findings of Bohr have given rise to
a very complex science which we will touch
upon briefly of quantum mechanics. And these
were based on a lot of the work done by Schrödinger
and also by Heisenberg. And the ideas of quantum
mechanics are thus: that energy only comes
in fixed quanta. In other words, you can’t
have an infinitesimal amount of energy. You
can only have energy which exists as certain
fixed packets – fixed packets of energy.
The other part of it – and this relates
specifically to electrons – is that everything
can be thought of as either a particle or
a wave. And the wave-like properties of a
large object, for example people and items
in the world around us, are very, very small.
However, the wave-like property of things
which are really, really small, like electrons,
is actually quite pronounced. And this was
where the original planetary model of electrons
orbiting around a nucleus was flawed because
it implied only that an electron had a particle-like
property rather than a wave-like property.
04:51
So this relates, as I said before, to the
uncertainty principle and wave equations for
electrons and other particles. And what these
mean is that, when you’re looking at an
electron, you can’t talk about an electron
as being in a discrete part of a shell. All
you can do is determine where is the greatest
probability of being able to find it. And
this may seem irrelevant in the context of,
let’s say, covalent and ionic bonding going
forward but it’s very important because
it relates to the existence, shape and energy
of orbitals as we will see a little later
on.
05:27
And so, when we’re looking at the application
of these principles and the application of
quantum mechanics and their associated equations
to an atom structure, we tend to consider
electrons as waves and therefore we’re looking
at the probability or probability density
of finding an electron at a specific distance
from a nucleus in a specific shape.