Which of the following best describes how the first ionization energy changes within a group and across a period in the periodic table?

• A. It increases down a group and decreases across a period from left to right.
• B. It increases both down a group and across a period from left to right.
• C. It decreases down a group and increases across a period from left to right.
• D. It remains constant both down a group and across a period.

Answer: (C)

Verified step by step guidance

1. Recall that ionization energy is the energy required to remove an electron from a gaseous atom or ion.
2. Understand the trend down a group: as you move down a group in the periodic table, the atomic radius increases because new electron shells are added. This increased distance and shielding effect make it easier to remove an electron, so the ionization energy decreases.
3. Understand the trend across a period: as you move from left to right across a period, the nuclear charge increases (more protons), but electrons are added to the same shell, so shielding remains relatively constant. This stronger attraction pulls electrons closer, making them harder to remove, so ionization energy increases.
4. Summarize the trends: ionization energy decreases down a group due to increased atomic size and shielding, and increases across a period due to increasing nuclear charge with constant shielding.
5. Therefore, the correct description is that ionization energy decreases down a group and increases across a period from left to right.