Why does the effective nuclear charge experienced by valence electrons decrease down the group?

The outermost electrons are farther away from the nucleus

The correct answer is option 4. The outermost electrons are farther away from the nucleus.

Here's why:

Effective nuclear charge \((Z_{eff})\) is the positive charge experienced by an electron in an atom, considering the partial shielding effect of inner electrons.

As you move down a group in the periodic table, the number of electron shells increases.

This means the valence electrons (the outermost electrons) are further away from the positively charged nucleus.

Due to the increased distance, the attraction between the nucleus and the valence electrons weakens.

The inner electrons partially shield the valence electrons from the full positive charge of the nucleus, but this effect also decreases with distance.

As a result, the effective nuclear charge experienced by the valence electrons decreases down the group.

Therefore, option 4 accurately explains why the effective nuclear charge decreases down the group.