Electron Configuration

Ah, electrons. Those tiny little subatomic particles, whizzing around in their orbits. They are almost 2000 times lighter than a proton and only a third of the diameter, but they are extraordinarily important. You'll remember from Fundamental Particles that whilst the number of protons tells you the element of the atom, the number of electrons and their configuration give you its reactivity and chemical properties. These are major roles for such tiny particles!

What is electron configuration?

Electron configuration is also known as electronic configuration. It is the arrangement of electrons in shells, sub-shells, and orbitals within the atom.

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 one 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.commons.wikimedia.org

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 orbital, and which shell and sub-shell they are found in.

How can we find electron configuration?

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.

A diagram showing how energy level varies between different sub-shells. The arrow shows an increase in energy. In general, as the principal quantum number n increases, energy levels increase. Note the exception to the rule, 3d.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.

To work out electron configuration, it can help to draw out the different orbitals and subshells. Each box represents an orbital and we use half arrows to show the electrons and their spin states. Remember that each orbital can contain a maximum of two electrons.

Let’s have a go at working out the electronic configurations in a few examples.

Carbon has a proton number of 6, meaning that it also contains six electrons. According to the Aufbau principle, electrons will fill the lowest energy level sub-shells first. Therefore, two electrons will first fill the single orbital in 1s. Two further electrons will then fill the single orbital in 2s, the sub-shell with the next lowest energy level. That leaves two electrons to go in 2p. However, according to Hund’s rule, the electrons will prefer to go into separate orbitals within a sub-shell. The overall electron configuration is shown below.

The electron configuration of carbon.StudySmarter Originals

Another example is magnesium.

Magnesium has twelve electrons. Like carbon, its first two electrons will fill 1s and the next two will fill 2s. The next six electrons will fill 2p, leaving two electrons remaining. These will both go in 3s, the next lowest energy level, as shown.

The electron configuration of magnesium.StudySmarter Originals

We can show how many electrons are in each sub-shell with simple notation. In this way, magnesium’s electron configuration is represented by , and carbon’s by . The superscript numbers show the number of electrons present.

You may have noticed a pattern. An element’s position on the periodic table relates to which sub-shell its outermost electron is 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 in. commons.wikimedia.org

Writing out the full electronic configuration of say, strontium, could get a little boring - after all, it has thirty eight electrons! If you look at the periodic table, you’ll see that strontium has just two more electrons than krypton. Strontium therefore has the same electron configuration as krypton, but also has two extra electrons. Using the above diagram, we can work out that they fill the 5s orbital. But instead of writing out all of the electrons in all of the sub-shells up to , we can simply put [Kr] . This shows that this element has the same configuration as krypton with two additional electrons in the 5s sub-shell. It can be done with any of the noble gases.

How are ions formed?

We know how to fill in sub-shells and orbitals with electrons, but how do they empty? Remember, ions are atoms that have gained or lost electrons.

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

Let’s look at an example.

Calcium atoms, Ca, have the electron configuration . However, ions, which have lost two electrons, have the configuration . This can also be written as .

Exceptions to the filling rules

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 . Hang on a second - why is there only one electron in the 4s sub-shell? We’d expect chromium's configuration to be ! 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 , not . This again is a slightly reduced energy arrangement due to the lack of electron-electron repulsion.