Ionization Energy (3) Sub Shell Atomic Structure and Hund's Rule

Ionization Energy (3) Subshell Structure of Atoms and Hund’s Rule.

In this post I’m going to discuss one of the more trickier topics I’ve found in many years of teaching that’s associated with ionisation energy. 

This post is about the evidence that exists for a more intricate structure of the electrons in an atom. 

It seems that using the ionization energies of an element to point to electron energy levels or shells is not all that we can work out from their values. 

If you thought you had reached the limits of discussion about shells with your school chemistry, well you would be wrong.

As I said in my previous post, looking at the typical representations of the electron energy levels, all levels except the first level are split into sub levels or sub shells. 

What is the justification for this sub shell structure of the electron shells in atoms?

Well, the evidence comes from a plot that you can make of the first ionisation energies of the elements.

Here is the plot of the values of E_m1 (kJ.mol^—1) against nuclear charge (Z) for the first twenty elements of the periodic table.

Several obvious things are clear from this plot. 

For example, the more protons in an atom the higher the first ionisation energy of that element. 

But that’s not absolutely true, it's only true with a shell of electrons if you look at the elements from lithium (Li) to Neon (Ne) which each have two electron shells the ionisation energy does increase. 

But as soon as you add a third shell the ionisation energy falls dramatically say compare Neon with sodium (Na)

Why is there this fall in ionisation energy from neon to sodium?

Clearly, sodium has three shells, as we have seen, and the third shell must be further from the nucleus than the second so the energy required to remove the first mole of electrons from a mole of sodium atoms is much lower than that required to do the same for a mole of neon atoms. 

But why is the increase in first ionisation energy from Lithium to Neon not regular?

Similar to the argument above for the difference in first ionisation energy between sodium and neon being because of an additional shell, so the drop in ionisation energy between beryllium (Be) and boron (B) is, it is argued, due to the existence of electron subshells as part of the second shell and all subsequent shells. 

It is argued that the second shell is divided into two groups of electrons, two in a subshell closer to the nucleus than the other six. 

Electron subshells are designated historically using the letters s, p, d and f. 

These letters originally related to the corresponding lines in the emission spectra of elements: so s stood for simple, p stood for principle, d stood for diffuse and f stood for fundamental lines in the spectra of different elements. 

With the second shell dividing into two subshells, the lower energy subshell was labelled 2s and the higher energy shell 2p. 

Adding in the numbers of electrons means we can now write the electron arrangement or configuration for Neon as 1s² 2s² 2p⁶

So why is there a break between nitrogen and oxygen?

Here are couple of arguments to start your thinking.

It’s nothing to do with another subshell, but something to do with half filled subshells. 

Nitrogen with three electrons in the 2p subshell has it half filled, the next electron would give us oxygen with an electron that’s easier to remove since its first ionisation energy is lower than that of nitrogen.

Why could the extra electron in oxygen be easier to remove?

One argument goes like this. 

That extra electron must be further from the nucleus than the others in oxygen for it to have a lower first ionisation energy. 

Another argument involves Hund’s Rule. 

Hund’s Rule was put together in the late 1920’s to explain this and other phenomena. 

It simply (?) states that electrons in atoms fill orbitals singly, spinning in parallel i.e. the same direction.

When further electrons are added, these make pairs in each orbital and spin in opposite directions.

Hund’s Rule introduces us to three new ideas.

First, electrons in shells spin either east to west or west to east.

Second, electrons exist within subshells in orbitals: a maximum of two per orbital.

Third, in a multi-electron sub shell e.g. 2p with a maximum of six electrons, the electrons fill their orbitals singly first then they pair up.

When they fill singly, the electrons always spin in the same direction.

In an orbital containing two electrons, the electrons are spinning in opposite directions. 

Let’s compare the electrons in boxes model for nitrogen and oxygen:

Here is the second row of the Periodic Table and it includes oxygen and nitrogen. 

The arrows indicate the direction of electron spin. 

Look at nitrogen, where the electrons in the 2p subshell are all spinning in the same direction (all the arrows point the same way) and there is only one electron in each “box” representing an electron orbital. 

But in oxygen, there is the additional electron in the 2p subshell and it is added to fill one of the three orbitals and it spins in the opposite direction (note the arrows)

So the additional electron could also be repelled from this orbital because two electrons have the same negative charge, and that would reduce oxygen’s first ionisation energy. 

There you have it: subshells and Hund’s Rule.

Incidentally, if you go to somewhere like Wikipedia for a definition of Hund’s Rule you will probably end up with the technical version. 

Your College and A level Chemistry courses probably don’t require that kind of definition at this level. 

But do check with your course content and/or course textbook to be absolutely sure.

In my next post, I’m going to be looking at how we can build electron configurations using these new sub shell and orbital ideas and linking all this stuff to the biggest crib sheet in the history of science—Mendeleev’s Periodic Table. 

Here’s a taster:

What is the auf bau principle and why does it look like this:—