Why electron affinity of fluorine less than that of chlorine?

The electron affinity refers to the energy change that occurs when an electron is added to a neutral atom to form a negative ion. In the case of fluorine (F) and chlorine (Cl), the electron affinity of fluorine is higher than that of chlorine. This means that fluorine has a greater tendency to gain an electron and form a negative ion compared to chlorine.

There are several factors that contribute to the difference in electron affinities between fluorine and chlorine:

Effective nuclear charge: The effective nuclear charge is the net positive charge experienced by an electron in the outermost energy level of an atom. Fluorine has a higher effective nuclear charge compared to chlorine because it has one less electron shell and a smaller atomic radius. The higher effective nuclear charge in fluorine attracts an additional electron more strongly, resulting in a higher electron affinity.

Electron-electron repulsion: Both fluorine and chlorine have seven valence electrons in their outermost energy level. However, in fluorine, these electrons are closer together due to its smaller atomic radius. This proximity leads to increased electron-electron repulsion, making it easier for fluorine to accept an additional electron and increase its electron affinity.

Electron configuration: The electron configuration of fluorine is 1s² 2s² 2p⁵, while chlorine has the configuration of 1s² 2s² 2p⁶ 3s² 3p⁵. Fluorine has only one vacancy in its 2p sublevel, which makes it energetically favorable for an electron to fill that vacancy and complete a stable, full valence shell. In chlorine, however, the 3p sublevel has three vacancies, making it less favorable for an additional electron to be added, resulting in a lower electron affinity compared to fluorine.

These factors collectively contribute to the lower electron affinity of chlorine compared to fluorine. It's important to note that while fluorine has a higher electron affinity than chlorine, both elements exhibit strong electron-attracting abilities compared to many other elements.