period, successive electrons are added to orbitals in the same principal quantum level and the shielding of the nuclear charge by the inner core of electrons does not increase very much to compensate for the increased attraction of the electron to the nucleus. Thus, across a period, increasing nuclear charge outweighs the shielding. Consequently, the outermost electrons are held more and more tightly and the ionization enthalpy increases across a period. As we go down a group, the outermost electron being increasingly farther from the nucleus, there is an increased shielding of the nuclear charge by the electrons in the inner levels.
In this case, increase in shielding outweighs the increasing nuclear charge and the removal of the outermost electron requires less energy down a group. From Fig. . (a), you will also notice that the first ionization enthalpy of boron ( Z = ) is slightly less than that of beryllium ( Z = ) even though the former has a greater nuclear charge.
When we consider the same principal quantum level, an s -electron is attracted to the nucleus more than a p -electron. In beryllium, the electron removed during the ionization is an s -electron whereas the electron removed during ionization of boron is a p -electron. The penetration of a s- electron to the nucleus is more than that of a p- electron; hence the p electron of boron is more shielded from the nucleus by the inner core of electrons than the s electrons of beryllium. Therefore, it is easier to remove the p -electron from boron compared to the removal of a s- electron from beryllium.
Thus, boron has a smaller first ionization enthalpy than beryllium. Another “anomaly” is the smaller first ionization enthalpy of oxygen compared to nitrogen. This arises because in the nitrogen atom, three p- electrons reside in different atomic orbitals (Hund’s rule) whereas in the oxygen atom, two of the four p- electrons must occupy the same p- orbital resulting in an increased electron-electron repulsion. Consequently, it is easier to remove the