Which statement correctly describes trends in atomic radii across a period and down a group, and the underlying cause?

• A. Across a period radius increases; down a group radius decreases.
• B. Across a period radius decreases due to increasing effective nuclear charge; down a group radius increases due to additional electron shells.
• C. Across a period radius decreases due to shielding effect; down a group radius increases due to decreased attraction.
• D. Across a period radius increases due to electron-electron repulsion; down a group radius decreases due to compacting.

Answer: (B)

The trend in atomic radii is governed by effective nuclear charge and shielding. As you move across a period, protons are added to the nucleus while electrons fill the same principal energy level. The inner shielding from core electrons doesn’t increase much, so the effective nuclear charge felt by the outer electrons rises. That stronger pull draws the outer electrons closer, making the atomic radius smaller.

Down a group, each step adds a new electron shell, placing outer electrons farther from the nucleus. The increased distance plus more shielding from the added shells reduces the nucleus’s pull on the outermost electrons, so the radius grows.

The best statement says: across a period, radius decreases due to increasing effective nuclear charge; down a group, radius increases due to additional electron shells. The other options mix up the causes—shielding and the direction of radius change across a period, or the effect of adding shells down a group.