Ionization energies are dependent upon theatomic radius. Since going from right to left on the periodic table, the atomic radius increases, and the ionization energy increases from left to right in the periods and up the groups. As we move down the group, size increases so the outermost electrons are very far away from the nucleus so the electrons are loosely bounded by the nucleus so it is easy to remove it. Moving across a period, atomic size decreases, so the outermost electrons are nearer to the nucleus. So more force of attraction holds the electrons so more energy is needed to remove the electrons.
Successive ionization enthalpy
Ionization energy is the energy required to remove an electron from the outermost shell of an isolated gaseous atom. When the first electron or the most loosely bound electron is removed, the amount of energy required is less than the energy required to remove the electron in the next successive shell. This ionization energy goes on increasing with the number of electrons removed. So the number of electrons removed from the successive no of shells and the energy involved is called successive ionization energy.
Shielding effect can be defined as a reduction in the effective nuclear charge on the electron cloud, due to a difference in the attraction forces of the electrons on the nucleus. It is also referred to as the screening effect or atomic shielding.
Trend in ionization energy down the group
Ionization energy decreases down the group due increase in the atomic size (addition of new shell).
ionization enthalpy (ionization potential)
The ionization energy (IE) is qualitatively defined as the amount of energy required to remove the most loosely bound electron, the valence electron, of an isolated gaseous atom to form a cation.