The ease with which atoms gain and lose electrons strongly influences their properties - the ease with which they undergo chemical reactions and the types of chemical bonds that they form.

The ionisation energy is the enthalpy change for the removal of the least tightly bound electron from an atom or an ion in the gaseous state. Ionisation energies are given per mole of atoms or ions of a given type.

The first ionisation energy is that required for the removal of one electron.

For example: Na(g)  ® Na⁺(g) + e^-      DH⁰ = 495.8 kJ

Such reactions are always endothermic and ionisation energies are always positive - energy must be put into the system to pull an electron away from the nuclear charge.

Removing electrons from ions of increasing positive charge is increasingly difficult.

Note that the ionisation energy is defined as the energy required for the removal of the least tightly bound electron.

Electrons always come off in the order of the principal energy levels; for example, n = 3 electrons before n = 2 electrons etc.

 

• Ionisation energies generally increase across the periods
• Ionisation energies generally decrease down the families
• Removal of electrons from filled and half-filled subshells requires higher energy
• In successive ionisation energies, a large increase occurs for removal of electrons from noble gas or pseudo-noble gas configurations
• Metals have low ionisation energies
• Nonmetals have high ionisation energies

Information provided by: http://www.bpx.cov.ac.uk