Nitrogen has an electronic configuration of 1s2 2s2 2p3, which features a half-filled p-subshell. According to Hund's rule and exchange energy principles, half-filled orbitals possess extra stability. Removing an electron from this stable configuration requires more energy compared to oxygen (1s2 2s2 2p4), where electron-electron repulsion in the p-orbital makes the removal of the fourth electron relatively easier.

The assertion is false because electron affinity generally decreases (becomes less negative) as you move down a group due to increased shielding and distance from the nucleus. The reason is true because atomic radii do increase down a group as additional electron shells are added. Therefore, the assertion is incorrect while the reason is a scientifically established fact.

In the third period, chlorine (Cl) has the highest electron affinity. Electron affinity generally increases across a period as the effective nuclear charge increases and the atomic radius decreases, making it easier for the atom to accept an additional electron to complete its valence shell. Chlorine's high value is consistent with its position as a halogen.

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.