Intermoleculer Forces

Introduction :

We have enough reasons to believe that a net attractive force operates between molecules of a gas. Though weak in nature, this force is ultimately responsible for liquefaction and solidification of gases. But we cannot explain the nature of this force from the ideas of ionic and covalent bond developed so far, particularly when we think of saturated molecules like H2, CH4, He etc. The existence of intermolecular attraction in gases was first recognised by Vanderwaal’s (1813) and accordingly intermolecular forces have been termed as Vanderwaal’s forces. It has been established that such forces are also present in the solid and liquid states of many substances. Collectively they have also been termed London forces since their nature was first explained by London using wave mechanics.

Nature of Vanderwaal’s Forces:  

The Vanderwaal’s forces are very weak in comparison to other chemical forces. In solid NH3 it amount to about 39 KJ mol–1 (bond energy N-H bond = 389 KJ mol–1). The forces are non directional. The strength of Vanderwaal’s force increases as the size of the units linked increases. When other factors (like H-bonding is absent), this can be appreciated by comparison of the melting or boiling points of similar compounds in a group.

Origin of Intermolecular Forces:

Intermolecular forces may have a wide variety of origin.

Dipole-dipole interaction: This forces would exist in any molecule having a permanent dipole e.g. HF, HCl, H2O etc.

Ion-dipole interaction: These interaction are operative in solvation and dissolution of ionic compounds in polar solvents.

Induced dipole interaction: These generate from the polarisation of a neutral molecule by a charge or ion.

Instantaneous dipole-induced dipole interaction: In non polar molecules dipoles may generate due to temporary fluctuations in electron density. These transient dipole can now induce dipole in neighbouring molecules producing a weak temporary interaction.

Also Read :

chemical-bonding
Hybridization
Maximum Covalency & Resonance
Deviation from ideal behaviour & FAJAN’S RULE
Role of φ ( ionic Potential )
Hydrogen Bonding

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