Which of the following best explains why ionization energy decreases as you move down a group in the periodic table?

• A. Electron shielding decreases, allowing the nucleus to attract electrons more strongly.
• B. The number of protons decreases, reducing the nuclear charge.
• C. The number of valence electrons increases, making it harder to remove an electron.
• D. The atomic radius increases, causing the outer electrons to be farther from the nucleus and less tightly held.

Answer: (D)

Verified step by step guidance

1. Understand that ionization energy is the energy required to remove an electron from an atom in its gaseous state.
2. Recognize that as you move down a group in the periodic table, atoms have more electron shells, which increases the atomic radius.
3. Know that increased atomic radius means the outermost electrons are farther from the nucleus, reducing the electrostatic attraction between the nucleus and these electrons.
4. Consider the effect of electron shielding, where inner electrons repel outer electrons, further decreasing the effective nuclear charge felt by the outermost electrons.
5. Conclude that because the outer electrons are less tightly held due to increased distance and shielding, the ionization energy decreases as you move down a group.