Which statement best explains the trend in atomic radius across a period and down a group?

• A. Across a period radius decreases due to increasing effective nuclear charge; down a group radius increases due to additional electron shells and shielding.
• B. Across a period radius increases due to shielding; down a group decreases due to nuclear attraction.
• C. Across period radius constant; down group constant.
• D. Across period radius decreases due to ionization energy; down group increases due to electron repulsion.

Answer: (A)

When thinking about atomic size, the balance between the nucleus’s pull on electrons and the shielding effect of inner electrons matters most. Across a period (left to right), protons increase while electrons occupy the same shell, so the effective nuclear charge felt by outer electrons rises. That stronger pull draws the electron cloud closer, decreasing the radius.

Down a group, each step adds a new electron shell. The outer electrons move farther from the nucleus, and inner shells shield the nucleus’s pull, making the outer electrons less tightly held and increasing the radius.

So the statement that radius decreases across a period due to increasing effective nuclear charge, and increases down a group due to additional electron shells and shielding, best explains the trend. Other options misattribute the trends, such as claiming shielding dominates across a period or that nuclear attraction or electron repulsion alone governs the change, which doesn’t align with the observed patterns.