AS-Level The Periodic Table

  • The first ionisation energy of an element is the energy required to remove one mole’s worth of electrons from one mole's worth of gaseous atoms. 

  • As a trend, across a period in the periodic table, first ionisation energy increases.  

  • As a trend, down a group in the periodic table, first ionisation energy decreases.

  • The stability of the ion formed when an atom loses an electron determines the energy required to remove the electron.

  • Elements in groups 3 and 6 do not follow the general trend for first ionisation energy. It is easier to remove an electron from a group 3 element compared to a group 2 element and from a group 6 element compared to a group 5 element in the same period.

QUICK NOTES

First Ionisation Energies

 

The first ionisation energy of an element concerns the amount of energy required to remove an outermost electron from an element. It is not realistically possible to measure the exact energy required to remove one electron from one atom. Instead, the total energy required to remove one mole’s worth of electrons from one mole’s worth of atoms is measured. The atoms must be in a gaseous state.

 

 

As you go across a period, first ionisation energy increases.

 

Moving across a period, the atomic number of the elements increases, meaning they have more protons and electrons. An increase in the number of protons results in a nucleus having a greater positive charge. Added electrons go into the same electron sub-shell, meaning the electrostatic attraction between the nucleus and the outermost electrons increases. As the attraction is greater, the electrons are held more tightly to the nucleus, so it is harder to remove them.

 

This is a general trend, for information on the exceptions to this trend, please see below.

 

 

As you go down a group, first ionisation energy decreases.

 

Moving down the group, the number of electron sub-shells increases. This means that the outermost electrons have more inner sub-shells between them and the nucleus. These extra inner sub-shells effectively dilute the positive charge from the nucleus that reaches the outermost electrons; this is called shielding. The inner electrons experience a greater positive charge than the outer electrons, so the outer electrons are not held as tightly to the nucleus – it is easier to remove them.

 

Exceptions!

 

 

The general trend across a period does have a couple of exceptions. When dealing with ionisation energies, it is the stability of the ion produced that dictates how easy or difficult it will be to remove an electron.


For example, sodium is a group 1 metal, so it has one electron in its outer shell. To remove one electron would leave us with a positive sodium ion that has a full outer shell – a stable ion.

 

Fluorine, however, is a group 7 halide that has seven electrons in its outer shell. To remove one electron would leave us with a positive ion that has only 6 electrons in its outer shell – a very unstable ion! 

 

The electron configurations of an element can give us a good indication of how easily it will lose an electron.

 

Beryllium has the electronic configuration 1s 2s

Boron has the electronic configuration 1s 2s 2p

 

 

Removing an electron from beryllium, would produce an ion with the configuration 1s 2s .  Removing an electron from boron, would produce an ion with the configuration 1s 2s .

 

 

 

1s 2s  is more stable than 1s 2s 2p . This is because a 2p-orbital is higher energy than a 2s-orbital. Due to electron distribution around an atom or ion, filled orbitals are more stable than part filled orbitals.

 

For the same reason, 1s 2s  is more stable than 1s 2s , so it takes more energy to remove an electron from beryllium than boron.  Beryllium has a higher first ionisation energy than boron.

 

Oxygen has a lower first ionisation energy than nitrogen for similar reasons.

 

Nitrogen has the electronic configuration 1s 2s 2p , so the ion produced by removing an electron would be 1s 2s 2p .  This configuration is less stable than 1s 2s 2p .

 

 

 

 

 

 

 

 

 

 

Oxygen has the electronic configuration 1s 2s 2p , so the ion produced by removing an electron would be 1s 2s 2p . This configuration is more stable than 1s 2s 2p .

 

 

Electrons have a negative charge and repel each other. There are three p-orbitals (each orbital can hold two electrons), so when all the p-orbitals are half filled, the electrons within can repel equally and arrange themselves to be as far away from one another as possible. If we add another electron though, the electrons cannot be equally as far away from one another. This leads to increased electron repulsion and results in a higher energy configuration.

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