Ionization Enthalpy is the energy required to remove the most loosely bound electron from an isolated atom in the gaseous state for one mole of an element. It is expressed in kJ mol–1 (kilojules per mole).
M(g) + IE → M+(g) + e–
As you move from left to right in the periodic table, there is a nearly regular increase in the magnitude of the ionization enthalpy of elements.
On moving down a group the magnitude of the ionization enthalpy indicates a regular decline. The ionization enthalpy of the first member of any group is the highest within that group and the ionization enthalpy of the last member in the same group is the least.
The variation in the magnitude of ionization enthalpy of elements in the periodic table is mainly dependent on the following factors:
Size of Atom
In small atoms, the electrons are tightly held whereas in large atoms the electron are less strongly held. Thus, the ionization enthalpy decreases as the size of the atom increases.
When an electron is removed from an atom, the effective nuclear charge (ratio of the number of charges on the nucleus to the number of electrons) increases. As a result the remaining electrons come closer to the nucleus and are held more tightly. The removal of a second electron, therefore, requires more energy. The second ionization enthalpy is more than the first ionization enthalpy.
Type of Orbital
Since the orbitals (s, p, d and f) have different shapes, the ionization enthalpy depends on the type of electrons removed. An electron in an s orbital is more tighly held as compared to an electron in a p orbital. It is because an s electron is nearer to the nucleus as compared to a p electron. Similarily a p-electron is more tightly held than a d-electron, and a d-electron is more tightly held than a f-electron. If all other factors are equal, the ionization enthalpies are in the order s > p > d > f.