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## Electron configuration

For people studying chemistry, and spectroscopy within physics, the electron configuration nomenclature will be familiar. For example, the noble gas neon, which is at number 10 in the periodic table, may be written as $1s^{2} \; 2s^{2} \; 2p^{6}$. Titanium, which is at number 22 in the periodic table may be written as $1s^{2} \; 2s^{2} \; 2p^{6} \; 3s^{2} \; 3p^{6} \; 3d^{2} \; 4s^{2}$. What does this mean, and where do the letters come from?

## Energy levels within atoms

It was the Danish physicist Nils Bohr who, in 1913, suggested that electrons can only orbit the nucleus of an atom in certain allowed orbits. Chemists tend to refer to these levels as orbitals or shells. I will explain in a future blog how he calculated what these allowed orbits were, but the number before the s and p in the electron configuration of neon refers to the electron energy level, or what is sometimes called the principal orbit. So, $1s$ means the electrons are in the first or lowest energy level, the $n=1$ energy level as physicists would call it.

## The origin of the s,p,d,f orbital labels

The origin of the letters s,p,d and f in the electron configuration goes back to the origins of spectroscopy. In the period 1872-1880, George Liveing and James Dewar published a series of papers on their observations of the spectral lines of alkali metals.

They classified the lines based on their visual appearance, and came up with the nomenclature s,p,d and f.Their meaning is

1. s – sharp
2. p – principle
3. d – diffuse
4. f – fine (or fundamental)

There are letters beyond this, which just follow the alphabet, so g,h,i etc.

## The superscript after the letter

What does the superscript after the letter refer to? E.g. $1s^{2}$ or $2p^{6}$. We now know from what I have said above that the $1 \text{ in } 1s^{2}$ refers to the electron(s) being in the first orbital, the n=1 energy level. And the $2 \text{ in } 2p^{6}$ tells us the electrons are in the n=2 energy level. But what do the superscripts 2 after the “s” and 6 after the “p” refer to?

These refer to the number of electrons in that particular orbital/shell. So, for neon, there are 2 electrons in the 1s shell, 2 electrons in the 2s shell, and 6 electrons in the 2p shell. This makes a total of $2+2+6=10$ electrons in total, as one would expect given that neon has an atomic number of 10. For titanium we have 2 electrons in the 1s shell, 2 electrons in the 2s shell, 6 electrons in the 2p shell, 2 electrons in the 3s shell, 6 electrons in the 3p shell, 2 electrons in the 3d shell, and finally 2 electrons in the 4s shell. This makes a total of $2+2+6+2+6+2+2=22$, which agrees with the atomic number of titanium.

Why are there only 2 electrons in any of the s-shells, but 6 in the p-shells? I will explain this in a future blog, but it is to do with something called the Pauli exclusion priciple, which states that no two electrons (or, to be more general, no two fermions, and an electron is a fermion) can exist in the same quantum state. I will explain in a future blog how this leads to there being a maximum of two electrons allowed in any s-shell, a maximum of six in any p-shell, etc.

## Visualising all of this

The easiest way to visualise the electron configuration of an element is, of course, with a diagram. Below are the diagrams for neon and titanium. One can see the electron structure, with each n-level (1,2,3 etc) shown separately. But, these diagrams do not show the s-shell and the p-shell in the n=2 level separately, they just show 8 electrons (2 in the s-shell and 6 in the p-shell) in the n=2 level. To get the full story, we also need to have the electron configuration written out, as I did in the first paragraph.