Final answer:
The shielding effect influences atomic radius by modifying the effective nuclear charge felt by electrons, leading to smaller atomic radii across a period and larger ones down a group.
Step-by-step explanation:
The shielding effect is related to atomic radius in that it affects the effective nuclear charge (Zeff) perceived by electrons in an atom. Core electrons, which are adept at shielding, decrease the attraction between outer electrons and the nucleus. As you move from left to right across a period, the atomic number (Z) increases by one, but the increase in shielding is marginal, leading to a stronger Zeff. This stronger Zeff causes the valence electrons to experience a greater pull towards the nucleus, resulting in a smaller atomic radius. Conversely, as you move down a column in the periodic table, the valence electron shell increases in size due to a higher principal quantum number, which makes the atomic radius larger despite the shielding effect.