Which statement best explains the trend in atomic radius as you move across a period from left to right on the periodic table?

• A. Atomic radius increases because additional electron shells are added.
• B. Atomic radius remains constant because the number of protons and electrons increases equally.
• C. Atomic radius increases because electrons are added to the same energy level, causing greater electron-electron repulsion.
• D. Atomic radius decreases because the effective nuclear charge increases, pulling electrons closer to the nucleus.

Answer: (D)

Verified step by step guidance

1. Understand that atomic radius refers to the size of an atom, typically measured from the nucleus to the outer boundary of the electron cloud.
2. Recognize that as you move across a period from left to right, electrons are added to the same principal energy level (same shell), not to new shells.
3. Note that the number of protons in the nucleus increases across a period, which increases the positive charge of the nucleus (nuclear charge).
4. Consider the concept of effective nuclear charge (\[Z_{eff}\]), which is the net positive charge experienced by electrons after accounting for shielding by inner electrons. As protons increase, \[Z_{eff}\] increases.
5. Conclude that the increased effective nuclear charge pulls the electrons closer to the nucleus, reducing the atomic radius despite the addition of electrons to the same shell.