The ionization energy generally, increases across the periodic table as: Increased nuclear charge: Moving across the period, the proton number increases by 1 and so greater nuclear charge increases the force of attraction between the electrons and the positively charged nucleus. Thus, more energy is required to remove the electron. Decreased atomic radii. The electron to be removed is closer to the nucleus, resulting in greater electrostatic force of attraction. However, there are slight decrease in ionization energy when we move from: 1. Group II and Group III: The electron shifted from s to p orbital of the same shell, which increases the distance between electron and nucleus, so reduced Ionization Energy. 2. Group V and Group VI: This is the result of ‘spin-pair repulsion’ between electrons of Group VI (due to pairing of electrons)