Electron configuration is a way of describing the arrangement of electrons in an atom or ion. It's an important concept in physical chemistry and helps us understand the reactivity and chemical properties of an element. But how do we figure out the electron configuration of an element or ion?
In this article, we'll explore the basics of electron configuration. We'll define it, look at how it's represented, and learn about the Aufbau principle and Hund's rule. We'll also work through some examples to help you understand how to calculate electron configuration. Finally, we'll look at the evidence that supports electron configuration.
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 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.
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.
Orbitals are like tiny spaces where electrons hang out. It's like a tiny electron clubhouse! Each orbital can only have two electrons and they have to have different spins. The shape of the orbital depends on which sub-shell it's in. So, electron configuration is really just a way of saying how many electrons are in each orbital, and which shell and sub-shell they're in.
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.
When it comes to filling up orbitals with electrons, they follow a pretty simple rule: they always go for the lowest energy level first. Just like how you'd rather sit in a comfy chair than a hard one, electrons prefer to be in a lower energy state. This means they start filling up the shells with the lowest principal quantum numbers first. Inside each shell, they start with the s sub-shell, then move on to the p sub-shell, and finally the d sub-shell. But there's one exception: the 3d sub-shell actually has a lower energy level than the 4s sub-shell, so it gets filled first. Here's a helpful diagram to show you the different energy levels of each sub-shell.
Electrons are like negative magnets - they repel each other, so they prefer to keep their distance. This means that within sub-shells, electrons like to have their own space and will fill up empty orbitals first. These two rules are the foundation of electron configuration. But before we can start figuring out the electron configurations of different elements, we need to learn how to represent them.
We have two different ways of representing electron configuration: Standard notation. Box form.
The first way to represent electron configuration is by using standard notation, which is the easiest method. You simply list the electron sub-shells and indicate the number of electrons they contain with a superscript number. However, empty sub-shells can be left out.
For example, let's take Carbon which has two electrons in each of the 1s, 2s and 2p sub-shells. Its electron configuration using standard notation would be: 1s2 2s2 2p2.
When it comes to heavier elements, writing out all the different sub-shells can get tiring. To simplify this, scientists use shorthand standard notation. If an element has the same electron configuration as a noble gas, with the addition of a few extra electrons, then the name of the noble gas is written in square brackets and the extra electron sub-shells are added as usual. For instance, Strontium has the same electron configuration as Krypton, but with two extra electrons in the s sub-shell. Therefore, its electron configuration using shorthand standard notation would be: [Kr] 5s2.
Box form is a more detailed way of representing electron configuration compared to standard notation. It shows the position of electrons within individual orbitals. You represent the different orbitals in each sub-shell using square boxes and use vertical arrows to show the electrons. Traditionally, the first electron in each orbital is drawn pointing up, and the second pointing down. Let's take a look at the electron configuration of carbon in box form:
1s ↑↓
2s ↑↓
2p ↑↓ ↑
To work this out, we first start with the sub-shell with the lowest energy, which is 1s. This sub-shell can hold a maximum of two electrons, so we draw two boxes, one pointing up and the other pointing down to represent the two electrons. We then move on to the next sub-shell, which is 2s. This also holds a maximum of two electrons, so we draw two more boxes, one pointing up and one pointing down. Finally, we move on to the 2p sub-shell, which can hold a maximum of six electrons. We draw three boxes, with two arrows pointing up and one pointing down to show the four electrons in the sub-shell. This is how we represent the electron configuration of carbon using box form.
To work out the electron configuration of carbon using box form, we use the Aufbau principle and Hund's rule. Carbon has six electrons, so the first two electrons fill the single orbital in 1s. The next two fill the single orbital in 2s, which is the sub-shell with the next lowest energy level. The remaining two electrons go in 2p, but Hund's rule states that electrons prefer to go into separate orbitals within a sub-shell. Therefore, the 2p sub-shell is filled as follows:
1s ↑↓
2s ↑↓
2p ↑ ↑
The electron configuration of sodium using standard notation is: 1s2 2s2 2p6 3s1. Sodium has 11 electrons, and the first two fill the 1s sub-shell. The next two fill the 2s sub-shell, and the remaining seven electrons fill the 2p and 3s sub-shells.
To represent the electron configuration of oxygen using box form, we start with its first two electrons filling the 1s sub-shell. The next two electrons fill the 2s sub-shell. The remaining four electrons fill the 2p sub-shell, but since there are only three orbitals available, one of the electrons must pair up with another electron in an already-occupied orbital. The electron configuration for oxygen in box form is:
1s ↑↓
2s ↑↓
2p ↑ ↓ ↑ ↑
Finally, an element's position on the periodic table relates to which sub-shell its outermost electron is in. For example, a neutral atom from group 2 always has its outer electron in an s sub-shell, and a transition metal has its outer electron in a d sub-shell.
When ions form, they can gain or lose electrons. When gaining electrons, Hund’s rule and the Aufbau principle are followed as usual, forming a negative anion. When losing electrons, electrons are lost from the highest energy level first, forming a positive cation. However, there is an exception to this rule - 4s electrons are lost before 3d electrons.
For example, let's look at 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, which in this case 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.
There are some exceptions to electron configuration, such as chromium and copper. Chromium, Cr, has 24 electrons and the configuration 1s2 2s2 2p6 3s2 3p6 4s1 3d5. This is because the 4s and 3d sub-shells are very similar in energy level, and the lone electron in 4s doesn’t experience any repulsion because it isn’t paired up. This reduced electron-electron repulsion makes up for the fact that there is an extra electron in the slightly higher 3d energy level. Copper, Cu, also has a slightly reduced energy arrangement due to the lack of electron-electron repulsion in the 4s sub-shell. Its configuration is 1s2 2s2 2p6 3s2 3p6 4s1 3d10. In summary, when ions form, electrons can be gained or lost. Exceptions to the electron configuration occur when the energy levels of sub-shells are similar, causing a slightly reduced energy arrangement.
In conclusion, electron configuration describes the arrangement of electrons in an atom and is determined by the Aufbau principle and Hund's rule. When ions form, electrons can be gained or lost, with exceptions to the usual filling rules in the case of the 4s and 3d sub-shells. Atomic emission spectra and ionisation energies provide evidence for electron configuration, with big jumps between successive ionisation energies indicating the loss of electrons from different electron shells. Understanding electron configuration is essential for understanding chemistry and the behaviour of atoms and molecules.
What is electron configuration?
Electron configuration is the arrangement of electrons in shells, sub-shells, and orbitals within the atom.
How do you work out electron configuration?
Electron configuration is worked out by filling the shells of lowest energy level first (Aufbau’s principle), and then by filling empty orbitals within each sub-shell first (Hund’s rule).
