The ionisation potential or ionisation energy is the energy required to remove the most loosely bound electron from an isolated gaseous atom to form a cation. The ionisation energy generally increases across a period on the periodic table and decreases down a group. This is mainly due to the increase in nuclear charge and the decrease in atomic radius as we move across a period. When moving down a group, the outer electrons are farther from the nucleus and more shielded by additional electron shells, leading to a decrease in ionisation energy.

Let's consider the given elements which are all in the third period of the periodic table: Na (Sodium), Mg (Magnesium), Al (Aluminum), and Si (Silicon).

• Na has an electronic configuration of $$1s^2 2s^2 2p^6 3s^1$$. The single electron in the 3s orbital is relatively easy to remove.
• Mg has an electronic configuration of $$1s^2 2s^2 2p^6 3s^2$$. The removal of one electron from a filled 3s orbital requires more energy than removing one from a partially filled orbital, as in sodium.
• Al has an electronic configuration of $$1s^2 2s^2 2p^6 3s^2 3p^1$$. Its ionisation energy is less than that of magnesium because removing one electron from the 3p orbital, which is higher in energy and less shielded than the 3s orbital, is easier.
• Si has an electronic configuration of $$1s^2 2s^2 2p^6 3s^2 3p^2$$. The increased effective nuclear charge and additional electron in the 3p orbital mean that more energy is required to remove an electron from silicon than from aluminum.

Therefore, the trend in the first ionisation energies of these elements increases from Na to Mg due to the addition of electrons to the same shell where effective nuclear charge increases and shielding is similar. However, a decrease occurs moving from Mg to Al because electrons start filling a new subshell with slightly higher energy and less effective shielding. Then, silicon, again experiences an increase due to an increase in effective nuclear charge with the filling of the same p-subshell.

This trend corresponds to Option A: Na < Mg > Al < Si. This means:

• Sodium has the lowest first ionisation energy.
• Magnesium has a higher first ionisation energy than sodium.
• Aluminum has a lower first ionisation energy than magnesium, due to the starting population of the p-orbital.
• Silicon has a higher first ionisation energy than aluminum, owing to its greater effective nuclear charge and half-filled p-subshell stability.