Electron Gain Enthalpy and Electronegativity — Definition

Definition

Imagine an atom as a tiny, bustling city. This city has a central government (the nucleus) and various residential areas (electron shells) where citizens (electrons) reside.

Electron Gain Enthalpy (EGE): Now, suppose a new citizen (an electron) wants to move into this city. The 'electron gain enthalpy' is like the energy change that happens when this new citizen successfully moves in and settles down.

If the city (atom) is very welcoming and attracts the new citizen strongly, energy will be released, making the process favorable. This release of energy is represented by a negative sign for EGE. The more negative the value, the more eagerly the atom accepts the electron.

For example, halogens like chlorine are very eager to accept an electron to complete their outer shell, so they have highly negative electron gain enthalpies. However, if the city is already full or doesn't want new citizens (like noble gases or atoms with stable half-filled/fully-filled orbitals), it might even require energy to force the new citizen in, resulting in a positive EGE value.

The first electron gain enthalpy is usually negative for many elements, but adding a second electron to an already negatively charged ion (like adding an electron to O⁻ to form O²⁻) always requires energy input due to electrostatic repulsion, making the second EGE positive.

Electronegativity (EN): Now, consider two cities (atoms) deciding to share resources (electrons) to form a partnership (a chemical bond). Electronegativity is like the 'pulling power' one city has over the shared resources.

If one city is more 'electronegative,' it means it has a stronger ability to pull the shared resources (electrons) closer to itself within that partnership. It's not about gaining a whole new citizen, but about influencing the location of shared citizens.

For instance, in a bond between hydrogen and fluorine (HF), fluorine is much more electronegative than hydrogen. This means fluorine pulls the shared electrons closer to itself, making its side of the bond slightly negative and hydrogen's side slightly positive.

This property is crucial for understanding how atoms bond and how molecules behave. It's a relative scale, meaning we compare the pulling power of one atom against another, rather than measuring an absolute energy value.