Electron Affinity or Electron gain enthalpy (E.A.)
The electron affinity of an element may be defined as amount of energy which is released when an extra electron enters the valence orbital of an isolated neutral atom to form a negative ion.
Atom(g) + Electron(g) ® Negative ion(g) + Energy
The greater the energy released in the process of taking up the extra electron, the greater will be the electron affinity. Thus, ionisation potential measures the tendency of an atom to change into a cation (M ® M+ + le-) whereas the electron affinity measures the tendency of an atom to change into anion (X + e- ® X-).
Successive Electron Affinities. As more than one electron can be introduced in an atom, it is called second electron affinity for the addition of two electrons and so on. The first E. A. of active non metals is positive (exothermic) while the second E. A. even for the formation of oxide or sulphide ion is negative (endothermic). For example, the overall E.A. for the formation of oxide or sulphide ions has been found to be endothermic to the extent of 640 and 390 kJ mol-1 respectively.
X-(g) + e- + energy ® X2-(g)
It is interesting to note that the electron affinity of elements having a d10 s2 configuration has been found to be negative. This is so due to the accommodation of the electron in the higher p-orbital (Zn = -87 kJ mol-1, Cd = -56 kJ mol-1).
Elements of group 17 possess high electron affinity. The reason for this is that by picking up an electron halogens attain the stable noble gas electronic configuration.
The electron affinity is expressed in kJ mol-1.
Change of Electron Affinity along a Group. On moving down a group, the size of atom increases significantly and hence, the effective nuclear attraction for the electron decreases. Consequently the atom will possess less tendency to attract additional electrons towards itself. It means that electron affinity would decrease as we move down a group. In case of halogens the decrease in electron affinity from chlorine to iodine is due to steady increase in atomic radii from chlorine to iodine.
On moving down a group the electron affinity decreases. Thus, the electron affinity of Cl should be less than F. But actually the electron affinity of F (320 kJ mol-1) is less than Cl (348 kJ mol-1). The reason for this is probably due to small size of fluorine atom. The addition of an extra electron produces high electron density which increases strong electron-electron repulsion. The repulsive forces between electrons results in low electron affinity.
Electron affinities of noble gases are zero. As these atoms possess ns2np6 configuration in their valence shells, these are stablest atoms and there are no chances for the addition of an extra electron. Thus, the electron affinities of noble gases are zero.
Electron affinities of beryllium and nitrogen are almost zero. This may be due to the extra stability of the completed 2s-orbital in beryllium and of the exactly half-filled p-orbital in nitrogen. As these are stable electronic configurations, they do not have tendency to accept electrons and therefore, the electron affinities for beryllium and nitrogen are zero.
Change of Electron Affinity along a Period. On moving across a period, the size of atoms decreases and nuclear charge increases. Both these factors favour an increase in force of attraction exerted by the nucleus on the electrons. Consequently, the atom will possess a greater tendency to attract the additional electron, i.e., its electronic affinity would increase as we move from left to right. Due to this reason electron affinities of non-metals are high whereas those of metals are low.
Of all the metals, the E.A. of gold is comparatively high (222.7 kJ mol-1).
This value may be attributed to the higher effective nuclear charge and poor shielding of the nucleus by d electrons.