Which of the following best explains the decrease in first ionization energy when moving from nitrogen (N) to oxygen (O) across the periodic table?

• A. Oxygen has a lower nuclear charge than nitrogen, resulting in weaker attraction for its electrons.
• B. Nitrogen has a larger atomic radius than oxygen, so its electrons are farther from the nucleus and easier to remove.
• C. Oxygen's electrons are in a higher principal energy level than nitrogen's, requiring less energy to remove.
• D. Electron-electron repulsion increases in oxygen due to paired electrons in the 2p orbital, making it easier to remove an electron.

Answer: (D)

Verified step by step guidance

1. Recall that first ionization energy is the energy required to remove the outermost electron from a neutral atom in the gas phase.
2. Understand that as you move from nitrogen (N) to oxygen (O) across the periodic table, the nuclear charge increases because oxygen has more protons than nitrogen.
3. Recognize that despite the increased nuclear charge, oxygen has paired electrons in one of its 2p orbitals, while nitrogen has half-filled 2p orbitals with unpaired electrons.
4. Know that electron-electron repulsion between paired electrons in the same orbital increases the energy of those electrons, making them easier to remove compared to unpaired electrons.
5. Conclude that the decrease in first ionization energy from nitrogen to oxygen is best explained by the increased electron-electron repulsion in oxygen's paired 2p electrons, which outweighs the effect of the increased nuclear charge.