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 arrangement give you its reactivity and chemical properties. These are major roles for such tiny particles! But how do we find out the electron configuration of an element or ion?
- This article is about electron configuration in physical chemistry.
- It builds on your knowledge of electron shells, sub-shells and orbitals.
- We'll start by defining electron configuration before looking at representing electron configuration.
- We'll then learn about the Aufbau principle and Hund's rule.
- After that, we'll put our new knowledge to the test with plenty of worked examples that show you the electron configuration of various elements and ions.
- Next, we'll discover the exceptions to the filling rules.
- We'll finish by briefly exploring the evidence for electron configuration.
Electron configuration definition
Electron configuration, also known as electronic configuration, 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 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.
Carbon has two electrons in each of the 1s, 2s and 2p sub-shells. Write out its electron configuration using standard notation.
This is quite simple. We write the names of the sub-shells in a line, and use superscript numbers to show how many electrons they contain. In this case, each of the three mentioned sub-shells has just two electrons: 1s2 2s2 2p2.
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.
Strontium has the same electron configuration as krypton, but with two further electrons in the s sub-shell. Use shorthand standard notation to represent its electron configuration.
Once again, this is very straightforward - all we have to do is write [Kr] 5s2.
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.
Here's the electron configuration of carbon (1s2 2s2 2p2) in box form:
Electron configuration of carbon using box form. StudySmarter Originals
We'll look at how we worked out this electron configuration next.
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.
Use the Aufbau principle and Hund's rule to work out the electron configuration of carbon in box form.
You'll notice that this is the example we gave earlier, but now we'll talk you through how to do it.
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. This 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.
Electron configuration of carbon using box form. StudySmarter Original
Another example is sodium.
Give the electron configuration of sodium using standard notation.
Sodium has eleven 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 one electron. This goes in 3s, the next lowest energy level, as shown:
1s2 2s2 2p6 3s1
Next up: oxygen.
Give the electron configuration of oxygen using box form.
Oxygen has eight electrons. Its first two electrons fill 1s, whilst its second two fill 2s. Its next four go in 2p. Thanks to Hund's rule, the first three of these four are found in separate orbitals. However, the 2p sub-shell only has three electron orbitals, so the fourth and final electron has to double up and share an already-occupied one:
Electron configuration of oxygen using box form. StudySmarter Originals
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.
Give the electron configuration of Ca2+ ions.
Calcium atoms, Ca, have the electron configuration 1s2 2s2 2p6 3s2 3p6 4s2. When losing electrons, they lose them from the highest energy level first. In this case, that is 4s. Ca2+ ions have lost two electrons and so have the electron configuration 1s2 2s2 2p6 3s2 3p6 4s0. This can also be written as simply 1s2 2s2 2p6 3s2 3p6.
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.
Electron Configuration - Key takeaways
- Electron configuration is also known as electronic configuration and describes the arrangement of electrons in an atom.
- Electrons fill shells according to their energy levels, as dictated by the Aufbau principle and Hund’s rule. Electrons fill sub-shells with a lower energy level first, and within each sub-shell prefer to occupy their own orbital.
- When forming ions, electrons are usually lost from the higher energy level sub-shell first.
- Exceptions to the filling rules stem from the fact that the 4s and 3d sub-shells are similar in energy level. Always remember that the 4s sub-shell fills before the 3d sub-shell.
- First and successive ionisation energies give us evidence for electron configuration.
References
- Attribution: File:Atomic orbital energy levels.svg: Richard Parsons (raster), Adrignola (vector) derivative work: MikeRun, CC BY-SA 3.0, via Wikimedia Commons
- User:DePiep, CC BY-SA 3.0(https://creativecommons.org/licenses/by-sa/3.0/) , via Wikimedia Commons
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