Electron Configuration & Ionisation Energy (AQA A Level Chemistry): Video

Hi, I'm Eleanor with 3 years of experience teaching chemistry, and this video is about electron configuration and ionisation energy. The two are covered together because ionisation energy data is the evidence for the electron configuration of an atom or ion.

An electron configuration describes where the electrons in an atom sit: which shell, which sub-shell and which orbital. Ionisation energy is the energy needed to remove those electrons, and its size depends on where they sit. That is the link between the two topics. Measuring ionisation energies tells you what the electron configuration is.

This video covers them in that order. First, how electrons are arranged in shells, sub-shells and orbitals, and how to write a configuration in full and in shorthand. Then what ionisation energy is, how it is defined, and the four factors that affect its size. Then successive ionisation energies, where the jumps in the data mark the boundaries between shells and confirm the configuration.

This is the arrangement that the ionisation energy data later confirms.

Electrons are arranged around the nucleus in principal quantum shells, numbered by the principal quantum number, n. The lower the number, the closer that shell sits to the nucleus. The higher the number, the greater its energy and the further out it sits. Each shell holds a fixed number of electrons: two in the first, eight in the second, eighteen in the third, and thirty-two in the fourth.

Each shell is split into sub-shells, given the letters s, p and d. Within a shell, their energy increases in the order s, then p, then d. Shell one only has an s sub-shell, shell two has an s and p sub-shell only and shell three has all three.

Sub-shells contain orbitals, and each orbital holds a maximum of two electrons. An s sub-shell has one orbital, a p sub-shell has three, and a d sub-shell has five.

Orbitals have specific shapes. The s orbitals are spherical. The p orbitals are dumbbell shaped, and the three of them point at right angles to each other along the x, y and z axes. The shape of the d orbitals is not required.

The electron configuration shows how electrons are arranged in the shells, sub-shells and orbitals of an atom. The sub-shells are filled in order of increasing energy, starting at 1 s. The main exception is that the 3 d sub-shell is slightly higher in energy than 4 s, so 4 s is filled first.

Configurations are written either in full or in shorthand. Potassium has nineteen electrons, so in full it is 1 s 2, 2 s 2, 2 p 6, 3 s 2, 3 p 6, 4 s 1. The shorthand version replaces the filled inner shells with the symbol of the nearest preceding noble gas. For potassium that is argon, which accounts for eighteen of its nineteen electrons, so the shorthand is argon in square brackets, 4 s 1.

Ions are formed when atoms lose or gain electrons. For the transition metals, the 4 s sub-shell fills before the 3 d, but electrons are lost from the 4 s first, because once the 3 d sub-shell contains electrons the 4 s is the higher in energy of the two.

So far this is the arrangement on its own. The evidence for it comes from ionisation energy.

Ionisation energy is the amount of energy required to remove one mole of electrons from one mole of gaseous atoms of an element, to form one mole of gaseous ions. It is measured under standard conditions, and the units are kilojoules per mole. The first ionisation energy is the one that forms 1+ ions.

Its size is affected by four things: the size of the nuclear charge, the distance between the outer electrons and the nucleus, the shielding effect of the inner electrons, and spin-pair repulsion.

All four of those depend on where the electron sits: how far out it is, how many inner electrons lie between it and the nucleus, and what it shares an orbital with. That is why ionisation energy can be used as evidence for the electron configuration.

Those four factors produce a recurring pattern across the periodic table. Across a period, ionisation energy increases: effective nuclear charge rises, the atomic radius falls, and shielding stays fairly constant, so the outer electron is held more tightly and takes more energy to remove. Down a group, it decreases: more shells means a larger radius and more shielding, and those outweigh the increased nuclear charge. This explains why the elements in group one become more reactive down the group; they lose their outer electron more easily.

Successive ionisation energies are where the two topics come together.

After the first electron is removed, the second ionisation energy removes an electron from the 1+ ion, the third removes one from the 2+ ion, and so on until only the nucleus is left. That sequence is the successive ionisation energies.

They always increase, because each electron is being removed from an increasingly positive ion, with less shielding and a higher proton-to-electron ratio. But the increase is not constant. It depends on the electron configuration, and that is what makes it useful.

On a graph of successive ionisation energies, the big jumps show a change of shell and the small jumps show a change of sub-shell. Reading where the first big jump falls tells you how many electrons are in the outer shell, which confirms the electron configuration and gives you the group the element belongs to. Sodium first jumps between the first and the second, so it is in group one. Magnesium first jumps between the second and the third, so it is in group two.

Sodium's configuration puts one electron in the outer shell, and the jump in the data falls in the same place. The configuration and the measurement agree, which is why the two topics are taught together.

There are two things examiners are strict about here.

The first is terminology. AQA mark schemes reject "p shell" and require "p sub-shell" or "p orbital", so always say sub-shell.

The second is state symbols. Equations representing ionisation energies must have gaseous state symbols on the atoms and the ions, but not on the electron. You will lose the mark if you leave them out, even when the question does not ask you to include them.

Electrons occupy shells, sub-shells and orbitals, and they fill the lowest energy sub-shell available first, which is why 4 s comes before 3 d. Ionisation energy measures how tightly the outer electron is held, so it increases across a period and decreases down a group. And because the jumps in successive ionisation energies mark the boundaries between shells and sub-shells, ionisation energy data is what confirms the electron configuration.

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