Electron affinity is energy released when an electron is added to an isolated atom in the gaseous state. For example,
X(g) + e- ------> X-(g) + Energy
(where X = F, energy 322 kJ mol; X = Cl, energy = 349 kJ mol; X = Br, energy = 324 kJ mol; X = I, energy = 295 kJ mol).
Halogens (elements of group 17) can take up an electron to acquire the stable noble gas configuration. Their values for electron affinity are thus very high. As we move across a period, electron affinity usually increases, and while going down a group it decreases. The electron affinity of an element is generally determined indirectly from thermodynamic data.
Although electron affinities of all the elements have not been determined, the following trends in electron affinities of some elements in the periodic table are evident.
Electron affinities generally decrease in moving down the group. This is expected on account of the increase in size of the atom on moving down the group. Due to increase in the size of atoms, the effective nuclear attraction for electrons decreases. As a result, there are fewer tendencies to attract additional electrons with an increase in atomic number.
In the figure electron affinities of halogens are plotted against their atomic numbers. It may be noted that contrary to expectation, the electron affinity of fluorine is lower than that of chlorine. This is because the fluorine atom has a very compact electronic shell due to its small size. The compactness of the fluorine shell results in electron-electron repulsion whenever an electron is introduced into its 2p shell. This is why its electron affinity is less than expected. In the chlorine atom, the 3p orbitals are not as compact as the 2p orbitals in fluorine atom. The chlorine atom, because of weaker electron-electron repulsion, more readily accepts the incoming electron. The electron affinity of chlorine is, therefore, higher than that of fluorine.
In the case of noble gases, the outer s and p orbitals are completely filled. No more electrons can be accommodated in these orbitals. Noble gases, therefore, show no tendency to accept electrons. Their electron affinities are zero.
Electron affinities generally increase as we move across a period from left to right. This is due to the increase in the nuclear charge, which results in greater attraction for electrons. In the second period, for example, the electron affinity has a maximum value for fluorine.
O-(g) + e- -----> O2-(g)
Since a negative ion (O-) and electron repel each other, energy is required and not released by the process. All second affinities are thus negative.