Chemistry Ionic Bond and Lattice Enthalpy

Topics Covered :

● Ionic or Electrocovalent Bond
● Lattice Enthalpy

Ionic or Electrovalent Bond :

`color{purple}(✓✓)color{purple} " DEFINITION ALERT"`
When a bond is formed by complete transference of electrons from one atom to another so as to complete their outermost orbits by acquiring 8 electrons or 2 electrons in case of hydrogen,lithium etc. and hence acquire the stable nearest noble gas configuration,the bond formed is called `color{red}("ionic bond or electrovalent bond")`.

`=>` 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 :

(i) The ease of formation of the positive and negative ions from the respective neutral atoms;

(ii) 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, `Δ_(eg)H`, 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.

`=>` 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.

`color{red}("Note :")` The ammonium ion, `NH_4^ +` (made up of two nonmetallic 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 in fig.

`=>` 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.

`color{red}("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.

`color{red}("Note :")` 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.

`color{purple}♣ color{Violet} " Just for Curious"`
`=>` The compounds usually exist in the solid state.
`=>` They exist as ions not as molecules.
`=>` Ionic compounds possess high melting point and boiling point.
`=>` They are soluble in solvents like water which are polar in nature and have high dielectric constant.
`=>` They are good conductors of electricity.

Lattice Enthalpy :

`color{red}("Definition :")` The Lattice Enthalpy of an ionic solid is defined as the energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions.

`color{red}("Example :")` The lattice enthalpy of `NaCl` is `788 kJ mol^(–1)`.

● This means that `788 kJ `of energy is required to separate one mole of solid `NaCl` into one mole of `Na^+ (g)` and one mole of `Cl^(–) (g)` to an infinite distance.

`=>` This process involves both the attractive forces between ions of opposite charges and the repulsive forces between ions of like charge.

`=>` The solid crystal being three dimensional; it is not possible to calculate lattice enthalpy directly from the interaction of forces of attraction and repulsion only. Factors associated with the crystal geometry have to be included.