Video Tutorial Electron Configurations

Quick Notes Electron Configurations

  • Electrons occupy orbitals of lowest energy around a nucleus.
    • Higher energy orbitals are only occupied if all lower energy orbitals are already full.
  • Electron configurations show how electrons are arranged around a nucleus within an atom or ion.
    • The first number shows the shell the electrons are in.
    • The letter shows the shape of the orbital the electrons are in.
    • The second number shows the number of electrons in the orbital(s).
  • Due to electron repulsion and orbital shapes, 4s orbitals are occupied before the 3d orbitals.

Full Notes Electron Configurations


Recap: Negatively charged electrons are attracted to the positive charge of a nucleus and want to be as close to it as possible. If the electrons get too close to each other they will repel, meaning electrons are unable to simply ‘crowd’ around a nucleus. Instead, they can only exist in certain regions of space – far enough away from each other to minimise repulsion, yet still as close to the nucleus as possible. These regions of space are called orbitals. One orbital can hold two electrons (one pair). For more detail, see Electron Orbitals.

As you move away from the nucleus of an atom, the amount of available space for electrons also increases; there are more possible orbitals that electrons can exist in. This is why as the shell number increases, the number of electrons that can be in the shell also increases.

The shapes of orbitals appear very strange as you move up to higher shells, this is only because there are more ways of arranging electrons around each other. At A-level, only the shapes of the s and p orbitals need to be remembered. The shapes of d-orbitals should be understood to help with transition metal ions.

Electron Configurations and the Aufbau Principle

Orbitals closest to a nucleus will contain electrons at a lower energy than the orbitals further away from it. Even if orbitals are in the same shell, they can still have different energies (called sub-shells). This means certain orbitals are filled before others.

Lower energy orbitals always fill with electrons before the higher energy orbitals. The order of this ‘filling’ is known as the Aufbau principle.

You should learn the types of orbitals that are present in each shell.

1st shell = s-orbital only
2nd shell = s-orbital and p-orbitals
3rd shell = s-orbital, p-orbitals and d-orbitals
4th shell = s-orbital, p-orbitals, d-orbitals and f-orbitals

How an atom’s electron orbitals are filled can help predict how it will react or the charged ions it may form. For this reason, electron configurations are used to show electron arrangement around an atom.

Take sodium, for example. Sodium has 11 electrons.

The lowest energy orbital available is 1s. Two electrons can occupy the s-orbital, meaning the 1s orbital contains two electrons. This is written as 1s2.

The next lowest energy orbital available is 2s. Again, two electrons can occupy this orbital, meaning the 2s orbital contains two electrons. This is written as 2s2.

The next lowest energy orbitals available are the 2p orbitals. These all have the same energy. There are three 2p orbitals, each orbital can hold 2 electrons, meaning the 2p orbitals contain six electrons. This is written as 2p6.

The next lowest energy orbital available is the 3s orbital. This orbital can fit 2 electrons in, but sodium only has 11 electrons in total and 10 electrons are already occupying lower energy orbitals, leaving only one 1 electron left over to go into the 3s orbital. This is written as 3s1.

Combing all these electron configurations together gives 1s22s22p63s1.

The electron in the higher energy 3s orbital is easily lost, forming a positive sodium ion that would have the electron configuration 1s22s22p6.

Electron configurations and 3d orbitals

3d orbitals are higher energy than a 4s orbital. This seems strange as surely the 3d orbitals are closer to the nucleus than the 4s orbital? We account for this because 3d orbitals have unusual shapes, leading to increased electron repulsion. This makes it easier to put electrons into the 4s orbital than the 3d orbitals. So, the 4s orbital is filled before the 3d orbitals.

What about d-block elements?

This is discussed in more detail with transition metals (see Transition Metals).

There is a major flaw with the above approach to 3d orbitals and 4s orbitals. If we are saying 4s orbitals are lower energy than 3d orbitals, surely transition metal elements would lose 3d electrons before 4s electrons when forming ions?

The problem is, that doesn’t happen!

The answer lies in the shapes of the 3d orbitals. If the 3d orbitals can spread their electrons out evenly, they actually become lower energy than the 4s orbital, meaning the 4s orbital electrons are higher energy and more easily lost. This results in some transition metals losing their 4s electrons before their 3d electrons.

It is always key to remember these are just useful models and simplifications can lead to ‘exceptions’.