From the Kössel and Lewis treatment of the formation of an ionic bond, it follows that the formation of ionic compounds would primarily depend upon:
- The ease of formation of the positive and negative ions from the respective neutral atoms;
- The arrangement of the positive and negative ions in the solid, that is, the lattice of the crystalline compound.
The formation of a positive ion involves ionization, i.e., removal of electron(s) from the neutral atom and that of the negative ion involves the addition of electron(s) to the neutral atom.
M(g) → M+(g) + e– ; Ionization enthalpy
X(g) + e– → X – (g) ; Electron gain enthalpy
M+(g) + X –(g) → MX(s)
The electron gain enthalpy, is the enthalpy change (Unit 3), when a gas phase atom in its ground state gains an electron. The electron gain process may be exothermic or endothermic. The ionization, on the other hand, is always endothermic. Electron affinity, is the negative of the energy change accompanying electron gain.
Obviously ionic bonds will be formed more easily between elements with comparatively low ionization enthalpies and elements with comparatively high negative value of electron gain enthalpy.
Most ionic compounds have cations derived from metallic elements and anions from non - metallic elements . The ammonium ion, NH4+ (made up of two non-metallic elements) is an exception. It forms the cation of a number of ionic compounds.
Ionic compounds in the crystalline state consist of orderly three - dimensional arrangements of cations and anions held together by coulombic interaction energies. These compounds crystallise in different crystal structures determined by the size of the ions, their packing arrangements and other factors.
The crystal structure of sodium chloride, NaCl (rock salt), for example is shown below.
In ionic solids, the sum of the electron gain enthalpy and the ionization enthalpy may be positive but still the crystal structure gets stabilized due to the energy released in the formation of the crystal lattice .
For example: the ionization enthalpy for Na+(g) formation from Na(g) is 495.8 kJ mol–1 ; while the electron gain enthalpy for the change Cl(g) + e– → Cl– (g) is, – 348.7 kJ mol–1 only.
The sum of the two, 147.1 kJ mol-1 is more than compensated for by the enthalpy of lattice formation of NaCl(s) (–788 kJ mol–1). Therefore, the energy released in the processes is more than the energy absorbed.
Thus a qualitative measure of the stability of an ionic compound is provided by its enthalpy of lattice formation and not simply by achieving octet of electrons around the ionic species in gaseous state.
Since lattice enthalpy plays a key role in the formation of ionic compounds, it is important that we learn more about it.