Electron arrangements (electronic configuration) and ionisation energies

This blog post is aimed at those of you who are at the very start of your IB course, and may not have even started your Diploma yet. It also serves as some useful revision for those of you who have completed your first year’s IB studies. If you are yet to start your Diploma course, I hope that it will allow you to get a heads up on things before your courses get underway.

The pre IB model for electron arrangements is usually taught that electrons are found in shells (that go around the nucleus).

We find 2e in the first shell, all other shells have 8e.

Eg,

Sodium, 11e → 2, 8, 1

Calcium 20e → 2, 8, 8, 2

Oxygen, 8e → 2, 6

But how do we know this? Where does the evidence come from?

The evidence comes from ionisation energies. So, what is ionisation energy?

The simple explanation is that it is the energy required to remove an electron from an atom or ion is the gaseous state.

However, the textbook adds to this, stating it is the amount of energy required to remove one mole of electrons from one mole of gaseous atoms or ions.

Eg, X (g) → X⁺ (g) + e^–

The units are therefore kJ mol ^-1

Some examples are show below, along with the reasons why they have been picked:

C (g) → C⁺ (g) + e^–             We wouldn’t expect to find C as a gas

Cl (g) → Cl⁺ (g) + e^–                We wouldn’t expect to find an atom of Cl or a + ion

As I said at the start, we will use ionisation energies to help explain electron arrangements, so let us start by looking at the evidence for shells:

We will consider the trend in first ionisation energy going down a group.

A graph of first ionisation energy v group 2 elements is shown below:

Group 2 elements

 

The trend that can be seen is that as you go down the group, the first ionisation energy decreases (in other words, as you go down the group, it gets easier to remove the outer shell electron).

To explain the reason for this, we need to be aware of three important concepts:

1. “Distance effect” – this is the electrostatic attraction between the negative electron and positive nucleus. This electrostatic attraction gets weaker as the distance increases.
2. “Shielding” – this is the effect that inner energy levels have on an outer shell electron. The inner shells shield the outermost negative electron from the positive nucleus, reducing the electrostatic attraction between the nucleus and the electron.
3. “Nuclear charge” – this is the number of protons in a nucleus.

Because the outer shell electrons get further away from the nucleus as we go down the group, this means that the electrostatic attraction between the outer shell electron and positive nucleus weakens as the electron is further away from the nucleus and there are more inner shells shielding the electron.

This is the first piece in the jigsaw that is used to provide evidence for shells. Not all electrons are in the same place in an atom. There are two other trends in ionisation energies that we use to develop this idea, successive ionisation energies and trends in ionisation energies across a period. More about this next month.