Electron Configuration

Electron configuration definition

If you’re not familiar with the above terms, we recommend looking at Electron Shells to learn a bit more about them. For now, we’ll just provide a quick summary.

Electron shells

Electron shells are also known as energy levels. Each shell has a specific principal quantum number. As shells get further from the nucleus, their principal quantum number increases and they have a higher energy level.

Electron sub-shells

Sub-shells are divisions within each shell. They also have different energy levels - the s sub-shell has the lowest energy, then p, then d, then f. Each sub-shell contains different numbers of orbitals. For example, the s sub-shell has just one orbital whilst p sub-shells have three and d sub-shells have five.

A graph showing the different energy levels of shells, subshells and orbitals, Richard Parsons (raster), Adrignola (vector) derivative work: MikeRun, CC BY-SA 3.0, commons.wikimedia.org[1]

Electron orbitals

Orbitals are regions of space where an electron can be found 95 percent of the time. Each orbital can contain at most two electrons. These electrons must have different spins - one has an up spin, the other a down spin. Orbitals also have different shapes depending on their subshell.

If we bring this all together, electron configuration is simply how many electrons are in each atomic orbital, and which shell and sub-shell they are found in.

Electron configuration rules

There are two main rules that you should know that will help you work out an atom’s electronic configuration. These are known as Hund’s rule and the Aufbau principle. We’ll take a look at both of them in turn before putting them into practice with some examples.

The Aufbau principle

First and foremost, electrons fill the sub-shell with the lowest energy level first. Atoms like being in a lower energy state and electrons are no different. In general, that means filling the shells with lower principal quantum numbers first, and within the shell first filling the s sub-shell, then the p sub-shell, then the d subshell. But remember the sneaky exception - 3d has a lower energy level than 4s! This means that it will be filled first. The diagram below reminds you of the energy levels of the different subshells.

The increasing energy of electron sub-shells. StudySmarter Originals

Hund’s rule

Electrons don’t really get along with each other. It makes sense - they are negative particles, and so if you put two of them close together, they will repel each other quite strongly. Because of this, within sub-shells electrons prefer to occupy their own orbital if they can, and so they will fill an empty orbital first.

These two rules form the basics of electron configuration. But before we have a go at working out the electron configurations of a few elements, we first need to learn how to represent electron configuration.

Representing electron configuration

We have two different ways of representing electron configuration:

• Standard notation.
• Box form.

Representing Electron Configuration: Standard notation

The first way of representing electron configuration is with standard notation. This is perhaps the easiest method; you simply list the electron sub-shells and indicate the number of electrons they contain with a superscript number. However, you don't need to worry about empty sub-shells - you can simply leave them out.

When representing the electron configurations of heavier elements, writing out all the different sub-shells get quite tiring. There's a way round this: if you know that a species has the same electron as a noble gas, with the addition of a few extra electrons, then you write the name of the noble gas in square brackets and add in the extra electron sub-shells as normal.

Representing Electron Configuration: Box form

Box form is a slightly longer way of representing electron configuration, but unlike standard notation, it shows the position of electrons within individual orbitals. You represent the different orbitals in each sub-shell using square boxes, and show electrons using vertical arrows. It is traditional to draw the first electron in each orbital pointing up, and the second pointing down.

Electron configuration of elements

We'll now put our new knowledge to the test with some examples. First, we'll work out the electron configurations of elements.

Another example is sodium.

Next up: oxygen.

You may have noticed a pattern. An element’s position on the periodic table relates to which sub-shell its outermost electron is found in. A neutral atom from group 2 always has its outer electron in an s sub-shell, for example, whilst a transition metal has its outer electron in a d sub-shell. This is shown below.

A diagram of the periodic table, showing how an element's position relates to which sub-shell its outer electron is found in.DePiep, CC BY-SA 3.0, commons.wikimedia.org[2]

Electron configuration of ions

We know how to fill in sub-shells and orbitals with electrons to form neutral atoms, but how do they gain or lose additional electrons to form ions?

• When gaining electrons, Hund’s rule and the Aufbau principle are followed as usual. This forms a negative anion.
• When losing electrons, electrons are lost from the highest energy level first - so in reverse order to filling up. This forms a positive cation. However, there is another sneaky exception to the rule: 4s electrons are lost before 3d electrons.

Let’s look at an example.

Exceptions to electron configuration

You’ll probably have guessed by now that although chemistry is a logical subject, there are always a few cases that seem to ignore all the standard rules. Unfortunately, you just have to learn them - although taking the time to understand why they misbehave can help you to remember them.

Take chromium. Chromium, Cr, has twenty four electrons and the configuration 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵. Hang on a second - why is there only one electron in the 4s sub-shell? We’d expect chromium's configuration to be 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁴ ! Well, this is because the 4s and 3d sub-shells are very similar in energy level. The lone electron in 4s doesn’t experience any repulsion because it isn’t paired up, and this reduced electron-electron repulsion makes up for the fact that there is an extra electron in the slightly higher 3d energy level. Atoms just like to be in the lowest energy state possible.

Likewise, copper, Cu, has the configuration 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰ , not 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁹. This again is a slightly reduced energy arrangement due to the lack of electron-electron repulsion.

A diagram showing the expected and observed configurations of chromium and copper. Note how both elements have just a single electron in 4s. This is because the lack of electron-electron repulsion creates a slightly lower energy arrangement. StudySmarter Originals

Evidence for electron configuration

To conclude this article, we'll briefly consider some of the evidence for electron configuration:

• Atomic emission spectra tell us that different quantum energy levels exist. Atomic emission spectra are produced when excited electrons emit light and return to their ground state, which is their lowest energy level. The wavelength and frequency of the light all depend on the electron's energy level.
• Successive ionisation energies also give us evidence for electron shells. Big jumps between an element's successive ionisation energies indicate that the electron is lost from a different electron shell that is closer to the nucleus.
• First ionisation energies give us evidence for sub-shells and orbitals. For example, the decrease in first ionisation energy between groups 2 and 3 shows that s and p sub-shells exist, whilst the decrease in first ionisation energy between groups 5 and 6 shows that the p sub-shell contains three orbitals.