Which of the following best explains the general trends and exceptions observed in ionization energy across the periodic table?

• A. Ionization energy generally increases across a period due to increasing nuclear charge, but exceptions occur when electron removal involves a new subshell or paired electrons, which can lower the required energy.
• B. Ionization energy decreases across a period because atomic radius decreases, making electrons easier to remove.
• C. Ionization energy remains constant across a period because the number of protons and electrons increases equally.
• D. Ionization energy increases down a group because atoms have more electrons, making them harder to remove.

Answer: (A)

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 the general trend that ionization energy increases across a period (left to right) on the periodic table because the nuclear charge (number of protons) increases, pulling electrons closer and making them harder to remove.
3. Note that ionization energy generally decreases down a group (top to bottom) because additional electron shells increase the distance between the nucleus and the outermost electrons, reducing the effective nuclear attraction.
4. Identify exceptions to the trend across a period, such as when removing an electron from a new subshell (e.g., from a filled s subshell to a p subshell) or from paired electrons in the same orbital, which can lower the ionization energy due to electron-electron repulsions or subshell stability.
5. Summarize that the best explanation combines the increase in nuclear charge across a period with the exceptions caused by subshell electron configurations and paired electrons, rather than simply atomic radius or electron count alone.