Q.Explain why the electron gain enthalpy of fluorine is less negative than that of chlorine.

The key idea here is the interplay between atomic size and inter-electronic repulsion affecting electron gain enthalpy.

Fluorine is significantly smaller than chlorine. When an electron is added to fluorine, it enters the compact 2p subshell. The existing electrons in this small volume exert strong inter-electronic repulsion on the incoming electron. This repulsion makes the addition of an electron less energetically favorable, resulting in a less negative (less exothermic) electron gain enthalpy. In contrast, chlorine is larger, and its valence electrons are spread over a larger 3p subshell. The incoming electron experiences much less inter-electronic repulsion, leading to …

Electron gain enthalpy ($\Delta_{eg}H$) is the energy change that occurs when an electron is added to a neutral gaseous atom to form a gaseous anion. A negative value indicates that energy is released (exothermic process), meaning the anion formed is more stable than the neutral atom. A more negative value signifies a greater tendency to accept an electron.

Generally, electron gain enthalpy becomes more negative across a period (due to increasing effective nuclear charge) and less negative down a group (due to increasing atomic size and shielding, which reduces the attraction for an incoming electron). However, there is a notable exception in Group 17 (halogens) when comparing fluorine and chlorine.

Let's break down why fluorine's electron gain enthalpy is less negative than chlorine's:

1. General Trend for Halogens: Following the general trend, as we move down Group 17 from fluorine to chlorine, bromine, and iodine, the atomic size increases. This increase in size typically means the incoming electron is added further away from the nucleus, experiencing weaker attraction and thus leading to a less negative (or less exothermic) electron gain enthalpy. So, we would expect fluorine to have the most negative electron gain enthalpy among the halogens.

2. The Anomaly: Fluorine vs. Chlorine: Contrary to the expected trend, chlorine actually has a more negative electron gain enthalpy than fluorine. The experimental values are approximately:

   • Fluorine (F): $-328 \text{ kJ mol}^{-1}$
   • Chlorine (Cl): $-349 \text{ kJ mol}^{-1}$ This indicates that chlorine has a greater tendency to accept an electron and forms a more stable anion than fluorine.

3. Atomic Size Difference: The key to understanding this anomaly lies in the significant difference in atomic size between fluorine and chlorine. Fluorine is the smallest element in its period and the smallest halogen. Its valence electrons are in the $2p$ subshell. Chlorine is larger, with its valence electrons in the $3p$ subshell.

4. High Electron Density in Fluorine: Due to its extremely small atomic radius ($64 \text{ pm}$), the $2p$ subshell of fluorine is very compact. The seven valence electrons are confined to a very small volume, leading to a high electron density around the nucleus.

5. Inter-electronic Repulsion in Fluorine: When an incoming electron approaches the fluorine atom to occupy the $2p$ subshell, it experiences strong electrostatic repulsion from the already densely packed electrons in that small $2p$ orbital. This repulsion significantly counteracts the attractive force from the nucleus. The energy released upon electron gain is thus reduced because some energy is required to overcome this strong electron-electron repulsion. …