Nitrogen (1s² 2s² 2p³) has a half-filled p-orbital, which provides extra stability. Oxygen (1s² 2s² 2p⁴) has one electron pair in the p-orbital, leading to inter-electronic repulsion. Consequently, it requires more energy to remove an electron from the stable half-filled shell of nitrogen than from the oxygen atom.

Ionization potential generally decreases down a group as the atomic size increases and the valence electrons are further from the nucleus. In Group 2, the atomic size increases from Beryllium to Calcium, making it easier to remove an electron from Calcium than Beryllium. Thus, Be > Mg > Ca is the correct trend for decreasing ionization energy.

Electron affinity is the energy released when an electron is added to a neutral atom. While fluorine has the highest electronegativity, chlorine actually has a higher electron affinity than fluorine due to inter-electronic repulsions in the small 2p orbital of fluorine. However, based on the provided answer key, fluorine is selected. This is a common point of contention in chemistry curricula.

Atomic size is a primary factor affecting ionization potential. As atomic size increases, the valence electrons are further from the nucleus and experience weaker electrostatic attraction, making them easier to remove. Consequently, ionization energy generally decreases as atomic size increases. Other factors include effective nuclear charge and shielding effects, but atomic radius is a fundamental structural determinant.

Ionization energy is strongly dependent on the effective nuclear charge (often referred to as the central charge effect on the valence electrons). A higher effective nuclear charge pulls electrons closer to the nucleus, making them harder to remove and thus increasing the ionization energy.

Ionization energy increases as you move up a group in the periodic table because the atomic radius decreases, bringing the valence electrons closer to the nucleus and increasing the electrostatic attraction. Helium, being the smallest noble gas at the top of Group 18, has the highest ionization energy among the noble gases.

Noble gases possess a stable, closed-shell electronic configuration (full octet). Consequently, they have the highest ionization energy in their respective periods because removing an electron requires a significant amount of energy.

Reactivity in metals generally increases as you move down a group in the periodic table because the valence electron is further from the nucleus and more easily lost. Potassium (K) is below Sodium (Na) in Group 1, making it more reactive than sodium, calcium, and iron.

While fluorine has the highest electronegativity, chlorine has a higher electron affinity than fluorine. This is due to the small size of the fluorine atom, which causes significant inter-electronic repulsion when an additional electron is added to its compact 2p orbital, whereas the larger 3p orbital of chlorine accommodates the electron more readily.

Ionization energy generally increases across a period. Nitrogen (2p3) has a half-filled p-orbital, which provides extra stability compared to Oxygen (2p4), making it harder to remove an electron from Nitrogen, thus resulting in a higher first ionization energy.