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 electron shells increases, causing electrons to be farther from the nucleus.
• B. Electrons are added to higher energy levels, increasing the size of the atom.
• C. The shielding effect increases significantly, allowing electrons to move further from the nucleus.
• D. The number of protons increases, resulting in a stronger effective nuclear charge that pulls 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), so the number of electron shells does not increase.
3. Note that the number of protons in the nucleus increases across a period, which increases the positive charge of the nucleus.
4. Understand that the increased positive charge (more protons) results in a stronger effective nuclear charge, which is the net positive charge experienced by electrons after accounting for shielding by inner electrons.
5. Conclude that this stronger effective nuclear charge pulls the electrons closer to the nucleus, decreasing the atomic radius across the period.