Electron gain enthalpy and electronegativity

Electron gain enthalpy, ∆egH, is the enthalpy change when an isolated gaseous atom gains an electron: X(g) + e⁻ → X⁻(g). If energy is released, ∆egH is negative: the atom 'likes' the extra electron. Halogens, one electron short of a noble gas configuration, have very negative values (chlorine –349 kJ mol⁻¹). Noble gases, with full shells, must take the electron into a new shell, so their values are positive (neon +116 kJ mol⁻¹).

The trend: ∆egH becomes more negative across a period and less negative down a group. But there is a surprise: fluorine (–328) is less negative than chlorine (–349), and oxygen (–141) less than sulphur (–200). The n = 2 shell is very small and crowded, so the incoming electron is repelled by the electrons already there.

Adding a second electron to a negative ion, such as O⁻ → O²⁻, means pushing it against repulsion, so the second electron gain enthalpy is positive.

Electronegativity is a different idea: the ability of an atom in a compound to attract the shared pair of electrons towards itself. It is not a measured energy, so it has no unit and cannot be measured directly. Pauling's scale sets fluorine at 4.0; oxygen is 3.5, nitrogen and chlorine 3.0, carbon 2.5.

Electronegativity increases across a period and decreases down a group, the same way as ionisation enthalpy. It is not a fixed value for an element: it changes a little with the element's state and the atoms it is bonded to.

Remember the difference: electron gain enthalpy is about an isolated atom; electronegativity is about an atom in a bond.