Ionic Bond
Ionic Bond
From the K ssel and Lewis treatment of the formation of an ionic
bond, it follows that the formation of ionic compound 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 compounds.
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.
The Electron gain enthalpy, ∆egH, is
the enthalpy change, 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 crystallize in different crystal
structures determine 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 form Na(g) is 495.8kJ
mol1; while the electron gain enthalpy for the change Cl(g) + e - Cl
(g) is, -348.7kJ mol1 only. The sum of the two, 147.1kJ mol1 is
more than compensated for by the enthalpy of lattice formation of NaCl(s) (-788kJmol1).
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.
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