Shells – Introduction to Chemistry

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.