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.