HELP on FIRST IONIZATION ENERGY eV (electron Volts)
The first ionization energy is the minimum energy required to remove an electron from an atom when that atom is in its gaseous state. Generally, it is the outermost electron that is removed first. First ionization energy varies periodically across the periodic table. In general it increases from left to right across a period, and decreases down a group. The increase across a period is as a result of a decrease in atomic size as more electrons are in the valence shell. This is because, the nearer the negatively charged electron to the positively charged nucleus, the more it is attracted, and the harder it is to wrench it from the atom. The decrease down a group is due to increasing atomic radius as more shells are added.
Note that when hydrogen, having just one electron, is ionized, it becomes a bare nucleus, in this case a single proton, (or deuteron, if it was heavy hydrogen). Also, when electrons are removed from an atom, the remaining electrons do not have the same energies as they had before, because the ion now has a net positive charge and the electrons experience a stronger attractive force to the nucleus.
First ionization energy is highest in the 'inert' noble gases, because they have a closed shell; and lowest in the alkali metals because they have a single s-electron in an outer shell and when that electron is removed, they then have the electronic structure of the noble gases, a closed shell.
The element with the highest first ionization energy is Helium at 24.6eV, that with the lowest is Caesium at 3.9eV.
