Which of the following best explains why atomic radius decreases as you move across a period from left to right in the periodic table?

• A. The number of protons increases, resulting in a greater effective nuclear charge that pulls electrons closer to the nucleus.
• B. Electrons are added to higher energy levels, increasing the size of the atom.
• C. The shielding effect increases significantly, reducing the attraction between the nucleus and outer electrons.
• D. The number of electron shells increases, causing the outer electrons to be farther from the nucleus.

Answer: (A)

Verified step by step guidance

1. Understand that atomic radius refers to the size of an atom, typically measured as the distance from the nucleus to the outermost electron cloud.
2. Recognize that as you move across a period from left to right, protons are added to the nucleus, increasing the positive charge (atomic number) of the nucleus.
3. Note that electrons are added to the same principal energy level (same shell) across a period, so the number of electron shells does not increase.
4. Consider the concept of effective nuclear charge (Z_eff), which is the net positive charge experienced by an electron after accounting for shielding by inner electrons. As the number of protons increases, Z_eff increases because shielding remains relatively constant within the same shell.
5. Conclude that the increased effective nuclear charge pulls the electrons closer to the nucleus, resulting in a decrease in atomic radius across the period.