The energy change that occurs when an electron is added to a neutral atom in the gaseous state to form a negative ion is called electron affinity. Because it measures the attraction, or affinity, of the atom for the added electron. For most atoms, energy is released when an electron is added.
The addition of an electron to a chlorine atom is accompanied by an energy change of -349 kJ/mol, the negativity sign indicating that energy is released during the process.
Thus, the electron affinity of chlorine is -349 kJ/mol.
Cl (g) + e- → Cl- (g) E.A. = -349 kJ/mol
Second Electron Affinity
Electron affinity defined above is strictly called first electron affinity. In addition to first electron affinity defined above, the second electron affinity of some elements like O, S and Se is also known.
The second electron affinity of an element, M(g) is defined as the amount of energy required to add one more electron to its mononegative anion, M(g)- to form a negative anion, M2- (g).
The addition of a second electron to uni negative ion is an endothermic process. This is due to the fact that the addition of an electron to a negative ion must overcome the repulsion due to the existing negative ion, hence the process is endothermic.
The formation of 02-(g) from 0(g) can be shown by the following two equations:
O(g) + e- → O(g) EA1 = -142 kJ/mol
O-(g) + e- → 02-(g) EA2 = +780 kJ/mol
Electron affinity measures the ease with which an atom gains an electron. Electron affinities similar to ionization potentials are expressed in electron volts or kJ/mol.
Note: Two sign conventions are used for electron affinity. In most introductory texts, including this one, the thermodynamic sign convention is used.
A negative sign indicates that the addition of an electron is an exothermic process, as in the
electron affinity given for chlorine, -349 kJ/mol.
Historically, however, electron affinity has been defined as the energy released when an electron is added to a gaseous atom or ion. Because 349 kJ/mol are released when an electron is added to Cl(g), the electron affinity, for this convention is +349 kJ/mol.
Large negative numbers such as for Cl -349 kJ/mol indicate that a very stable negative ion is formed. Small negative numbers indicate that a less stable ion is formed. The greater the attraction between a given atom and an added electron, the more negative the atom’s electron affinity will be. The electron affinity of Cl is the most negative of all the elements.
The electron affinity of an element depends mainly upon the following factors:
- Nuclear Charge. More the nuclear charge of the atom more strongly will it attract an additional electron. Therefore, electron affinity increases as the nuclear charge increases.
- Atomic size. The smaller the size of atom smaller will be the distance between the extra electron and the nucleus. Therefore, the electrostatic force of attraction will be more and the electron affinity will be higher.
- Electronic Configuration. Atoms having a stable electronic configuration (i.e., those having completely filled or exactly half-filled outer orbitals) do not show much tendency to add extra electron. Therefore, have either zero or very low, electron affinities.
In general, EA value decreases with the increasing atomic radius and increases with the decreased screening effect by the inner electrons.
Periodic Trends in Electron Affinities (Variation in electron affinity) in a period and in a group.
The electron affinity values generally increase on moving from left to right in a period in the periodic table to decrease in atomic size with the increase of nuclear charge. And thus, the nucleus would have increased attraction for additional electron.
The elements of Group ITA (Be, Mg, Ca—), Group VA (N, P, As—), and noble gases are exceptions.
The halogens, which are one electron shy of a filled p-subshell, have the most negative electron affinities. By gaining an electron, a halogen atom forms a stable negative ion that has a noble gas configuration.
Electron affinity values (kJ/ mol) of Some Elements
The exception mentioned above can be readily explained. The electron affinity values of Group llA metals are positive because they have already completely filled n-s orbitals. And the added electron will have to go to the n-p orbital of higher energy. Instead of releasing energy, the atom would absorb energy.
Be (1s2 2s2)+e- → Be- (s2 2s2 2p1); E.A > O
None of the Group llA forms stable negative ions. The addition of an electron to a noble gas would require that the electron resides in a new, higher-energy subshell. Occupying a higher-energy subshell is energetically unfavourable. So, the electron affinity is positive, meaning that the ion will not form.
Noble gases do not form stable anions. The values for Group VA elements are less negative (i.e. less exothermic) than expected. This is because, these elements have half-filled p-subshells. The added electron must be put in an orbital that is already occupied, resulting in larger electron repulsions.
As a result, these elements have electron affinities that are either positive (N) or less negative than the preceding Group IVA element.
Going down a group electron affinities undergo a general decrease. This is a result of the fact that the valence shell is progressively farther from the nucleus. With progressively more inner shells screening the nuclear charge. One important exception to this general trend is the low electron affinity of some second row (period) elements.
This effect is more evident in N, O, F, which all have substantially lower electron affinities than the corresponding third-period elements. The small size of the second-period elements causes serious electron-electron repulsions in the anion that forms when the atom accepts an electron. These repulsions lead to the lower electron affinities. It is difficult to determine the electron affinity as compared to ionization potential. Electron affinities can be obtained by using Born-Haber cycle.