What trend is observed in first ionization energy as you move down a group in the periodic table?

As you move down a group in the periodic table, the first ionization energy generally decreases. This trend is primarily due to two factors: increased atomic size and electron shielding.

As you move down a group, the number of electron shells increases. Each new shell is further away from the nucleus, which results in a larger atomic radius. The increased distance means that the outermost electrons are less strongly attracted to the positively charged nucleus.

Additionally, the effect of electron shielding becomes more pronounced. Electrons in the inner shells repel the outermost electrons, reducing the effective nuclear charge felt by these outer electrons. This makes it easier to remove an electron, hence requiring less energy to achieve ionization.

Consequently, as one goes down a group, the first ionization energy decreases, making it easier to remove an electron from these larger, more shielded atoms.