Electron configuration

Electron Configuration

‘Electrons-in-a-Box’ Notation

Orbitals and sub-shells

Orbitals are represented by boxes, electrons by an arrow. When 2 electrons occupy the same orbital, they are represented by arrows drawn in opposite directions (representing spin). In the graphic above, energy level/shell 1 has one sub-shell (s) with 2 electrons. Energy level 2 has 2 sub-shells (s & p). The p sub-shell for nitrogen has 3 orbitals with an electron in each.

Rules:

1. The lowest energy sub-shells fill with electrons first (Aufbau Principle)

Start with the 1s sub-shell first, then the 2s, then the 2p and so on.

Of course, there’s got to be an exception: The 4s sub-shell, important when studying Transition Metals, fills up before the 3d sub-shell. It turns out that the 4s sub-shell is lower in energy than the 3d sub-shell.

2. Each orbital is filled with one electron before they pair/share with another electron (Hund’s Rule)

Having two electrons in an orbital needs energy to overcome the repulsion between the electrons. Notice in the electron configuration for nitrogen, above, each 2p sub-shell contains only 1 electron.

3. Only two electrons, each having opposite spins, can fit in an orbital (Pauli Exclusion Principle)

Let’s look at calcium and figure out its electron configuration based on these rules…

The second shell has 2 sub-shells (2s & 2p). We have plenty of electrons left so we can completely fill these giving…

10 electrons have been used so far leaving 10 electrons we can use.

The third energy level also has 2 sub-shells (3s & 3p). The s subshell can take 2 electrons and the 3 orbitals of the 3p subshell can hold 6 – the configuration now reads…

The final 2 electrons now go into the 4s sub-shell giving the final configuration:

There is a gradual build-up of electrons starting from the lowest energy level. Here’s what that looks like for the first ten elements in the periodic table, from hydrogen to neon:

Ionisation energy

Calcium has 20 protons in the nucleus with 20 electrons. Remove an electron and a calcium ion is produced, with an overall positive charge…

Ca → Ca+ + e–

To remove this electron requires energy - the first ionisation energy…

The energy needed to remove one mole of electrons from one mole of gaseous atoms to form one mole of one plus ions in their gaseous state.

So, for a generalised metal, M: M(g) → M+(g) + e–

Successive electrons can be removed leading to second, third, fourth, fifth etc ionisation energies.

For the second ionisation energy…

The energy required to remove one mole of electrons from one mole of one plus ions in their gaseous state to form one mole of two plus ions in their gaseous state

Ca+ → Ca2+ + e–

There are 3 things that will affect the size of the ionisation energy:

1. Distance of an outer electron from the nucleus

In atoms, there is an attraction between the positive protons in the nucleus and the electrons. For an electron in an outer electron shell, the attraction lessens the further it is from the nucleus (the greater the atomic radius).

2. Nuclear charge

If there are more protons, the greater the force of attraction between electrons and the nucleus.

3. Electron shielding

This due to the repulsion between the outer electron and the electrons held in the inner shells. The repulsion lessens the attraction between the outer electrons and the nucleus.

Successive Ionisation Energies

Here are the successive ionisation energies for nitrogen.

There’s a gradual increase in ionisation energy. For successive ionisations, the remaining electrons in the outer shell are pulled slightly closer to the nucleus – there’s a greater attraction between the outer electrons and the nucleus. Consequently, the ionisation energy gradually increases.

Once the 5 electrons in the outer shell are removed, the ionisation energy significantly increases. This is because all electrons in the outer shell have been removed and the 6th electron is removed from the next shell. This shell is closer to the nucleus – the electrons ‘feel’ the pull of the nucleus more and are subject to less repulsion (nucleus is less shielded). More energy is needed to ionise the next electron.

This massive increase in ionisation energy also gives us a clue to the element. Since it occurs for the removal of the 6th electron, the atom must have 5 electrons in the outer shell so must belong to Group 5 of the Periodic Table.

In Period 3, the first ionisation energy generally increases from left to right as atomic radius decreases - it becomes more difficult to remove an electron as the electron is held more tightly by the nucleus.

However, there are certain exceptions to this trend, such as the significantly lower first ionization energy of aluminum compared to that of magnesium. These anomalies can be explained by considering the electron configuration of each element.

Magnesium: 1s2 2s2 2p6 3s2

Aluminium: 1s2 2s2 2p6 3s2 3p1

Aluminium's outer electron is in a p orbital and is shielded by the other electrons in the sub-shell. This p sub-level is of higher energy than the s sub-level and so less energy is required to remove this electron.

Phosphorous: 1s2 2s2 2p6 3s2 3p3

Sulfur: 1s2 2s2 2p6 3s2 3p4

A p sub-level has 3 sub-shells which can hold 2 electrons in each sub-shell. In phosphorous, the electrons are unpaired with parallel spins. Sulfur has 4 electrons so one of them must pair in a sub shell. As electrons are both negative particles, there is electron-pair repulsion.and it's easier to remove the unpaired electron in phosphorous - less energy is required.

Ionisation energy generally decreases down group 2. This is because the atomic radius of the elements increases, which reduces the effective nuclear charge (the attraction between the positively charged nucleus and the negatively charged electrons). The outermost electrons are therefore held less tightly and require less energy to remove.

The second ionization energy, which is the energy required to remove a second electron from a singly charged ion, is significantly higher than the first ionization energy for all elements, due to the increased attraction between the remaining electrons and the now positively charged ion.

The specific pattern of successive ionisation energies provides evidence for the electron configuration in different sub-shells and shells.