Across the period, the nuclear charge increases due to increase in proton number. However, shielding effect remains relatively constant as the extra electrons are added to the outermost quantum shell i.e. innermost electrons (determines the shielding effect) are not affected.
As a result, the effective nuclear charge increases which results in greater electrostatic attraction between the nucleus and outermost electron. More energy is required to remove the outermost electron. Hence, I.E. increases across the period.
Note: This is the reason why group I elements have the lowest I.E. while noble gases have the highest I.E..
Shortcut: In general, elements/ions having the valence electronic configuration ns1 will have the lowest I.E. while ns2 np6 will have the highest I.E..
Period 2 elements
Lowest – Li: 2s1
Highest – Ne: 2s2 2p6
Lowest – Be+: 2s1
Highest – Na+: 2s2 2p6
Lowest – B2+: 2s1
Highest – Mg2+: 2s2 2p6
Note: Therefore, we can observe a trend for each successive I.E., there is a shift to the right of the periodic table by 1 element.
Down the group, there is an increase in nuclear charge due to increase in proton number. Number of quantum shells also increases which increases the shielding effect. These 2 are opposing factors hence it is hard to determine which factor is more significant. Thus, we DO NOT talk about nuclear charge or shielding effect here at all.
Instead, we only discuss the distance of the valence electron from the nucleus. Due to an increase in quantum shells, the distance between the nucleus and valence electrons increases as well, which results in lesser electrostatic attraction between the nucleus and outermost electron. Less energy is required to remove the outermost electron. Hence, I.E. decreases down the group.
Next, we will examine some of the anomalies in ionisation energies.