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

• A. Electrons are added to higher energy levels farther from the nucleus.
• B. Shielding effect increases significantly across a period.
• C. Addition of electron shells increases the size of the atom.
• D. Increasing nuclear charge 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), not to higher energy levels.
3. Note that the number of protons in the nucleus (nuclear charge) increases across a period, which means the positive charge attracting the electrons becomes stronger.
4. Consider the shielding effect, which is the repulsion between inner electrons and outer electrons; across a period, shielding remains relatively constant because electrons are added to the same shell.
5. Conclude that the increasing nuclear charge pulls the electrons closer to the nucleus, reducing the atomic radius despite the addition of electrons.