Atomic orbitals, sub-orbitals and electron configurations

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At GCSE we learnt that electrons are arranged in shells or energy levels around a central nucleus – the first shell held up to 2 electrons, the second and third shells held up to 8 electrons. The atom looked like the solar system and this model was perfectly sufficient to explain what we needed to understand at GSCE.

Applying quantum mechanical ideas to Bohr’s model of the atom leads us to realise that not only are the energy levels surrounding the nucleus of an atom discrete or fixed or quantised, but that each energy level consists of different orbitals and sub-orbitals. Time to upgrade our working atomic model…

Orbitals and sub-orbitals

Each energy level holds electrons in orbitals and sub-orbitals. The further the energy level is from the nucleus, the more orbitals and sub-orbitals there are and the more electrons that energy level can hold (we tend not to describe energy levels as shells now).

The first energy level, closest to the nucleus, is called n=1 and consists of an s-orbital.

  • the spherical shape of an s-orbital describes the mathematical probability of finding an electron in this space surrounding the nucleus
  • an s-orbital can hold a maximum of 2 electrons, and the electrons are spin-paired to minimise repulsion between them (the easiest way to visualise this is to imagine one electron spinning clockwise on its axis and one spinning anti-clockwise)

The second energy level, n=2, consists of an s-orbital and p-orbitals, holding 8 electrons in total.

  • there are three p sub-orbitals, px, py and pz, each of which can hold 2 electrons
  • once again, the shape describes the mathematical probability of finding an electron in this space

The third energy level, n=3, consists of an s-orbital, p-orbitals and d-orbitals.

  • there are five d orbitals, each of which can hold 2 electrons

But just to make life a little more complicated, the energy of each of these orbitals begins to overlap as we move further from the nucleus. The diagram below shows this:

We can use this to write an electron configuration for an element based on a couple of rules (known as the aufbau principle) …

  1. orbitals are filled from the lowest energy up i.e. 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p
  1. the three p sub-orbitals are degenerate in atoms (they are of equal energy) so are filled individually before spin pairing electrons – the same applies to filling d orbitals. Hund’s rule states that the lowest possible energy configuration has the maximum number of parallel spin electrons. Parallel spin electrons have a lower electron-electron repulsion than spin paired electrons in the same sub-orbital.

There are two elements that have unexpected electron configurations that you need to know about – chromium and copper.

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Following a similar logic, the electron configuration for copper is 1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s1.

Practice questions

  1. Write out the electronic configuration for aluminium, iron, arsenic and bromine.

We often abbreviate the electronic configuration by starting with the closest noble gas and then adding the configuration of the outermost shells. For example, magnesium has the configuration 1s2, 2s2, 2p6, 3s2 which is the same as saying [Ne] 3s2 because neon has the configuration 1s2, 2s2, 2p6.

  1. Name the elements with the following electronic configurations:

(a) [He] 2s2 2p1

(b) [Ne] 3s2 3p2

(c) [Ar] 3d3 4s2

  1. Write out the electronic configuration using the noble gas shorthand notation for fluorine, sulfur and germanium.

Answers

  1. Al: 1s2 2s2 2p6 3s2 3p1 ; Fe: 1s2 2s2 2p6 3s2 3p6 3d6 4s2 ; As: 1s2 2s2 2p6 3s, 3p6 3d10 4s2 4p3 ; Br: 1s2 2s2 2p6 3s, 3p6 3d10 4s2 4p5
  2. (a) boron (b) silicon (c) vanadium
  3. fluorine: [He] 2s2 2p5 ; sulfur: [Ne] 3s2 3p4 ; germanium: [Ar] 3d10 4s2 4p2