Intermolecular forces are the forces of attraction between neighbouring molecules, situated at a distance much closer in comparison with their molecular diameter.
These forces arise due to the interaction between molecules. These forces are collectively called as van der Waal’s forces.
These forces due to the interaction between molecules are further classified as a) Dipole – dipole interaction b) Dipole – induced dipole interaction c) Dispersion forces or London Forces
The melting and boiling points of the substances depend on the magnitude of intermolecular forces. Larger the magnitude of these forces higher is the melting point and boiling point of the substance.
Besides van der Waal’s forces, there is one more intermolecular attraction force called as hydrogen bonding.
Dipole-Dipole Interaction Between Molecules:
Dipole – dipole interaction results in a force of attraction between neighbouring molecules having a permanent dipole moment. This interaction is due to electrostatic force.
Consider polar molecule HCl, a compound of electropositive element hydrogen and electronegative element chlorine. Due to which chlorine pulls shared electron in the bond towards itself acquiring small negative charge (δ–) and hydrogen acquires equal positive charge (δ+). Thus the HCl molecule becomes an electrical dipole.
Due to these, all molecules align in a way that oppositely charged ends come closer to one another as shown.
The strength of dipole-dipole interaction depends upon dipole moment of the interacting molecule.
Other examples of this type interaction are HBr, H2S, NH3.
This effect was studied by Keesom in 1912, hence these forces are also called Keesom forces and also referred as orientation effect.
Besides dipole-dipole interaction, the polar molecules interact with London forces. The combined effect of two increases the total intermolecular forces.
In dipole-dipole interactions, the interaction energy is inversely to the sixth power of the distance between the two interacting molecules.
Dipole – Induced Dipole Interaction Between Molecules:
This interaction was studied by Debye (1920). This interaction is also called as an inductive effect.
These type of forces operate between polar (μ > 1) and nonpolar (μ = 1) molecules.
In such interaction, the permanent dipole of the polar molecules induces dipole on the nonpolar molecule by deforming or polarizing its electronic cloud.
The interaction energy of these forces is inversely proportional to the sixth power of the distance between the two interacting molecules.
The strength of the forces depends on the distance between the molecules and polarizability of nonpolar molecule.
Example: Interaction between NH3 (polar) and C6H6 (nonpolar)
Ion – Dipole Interaction Between Molecules:
Any ion due to charge on it can interact with the oppositely charged site of polar molecule to develop the force of attraction between the ion and the polar molecule.
The strength of interaction depends upon the magnitude of the dipole moment of the polar molecule, size of the molecule, charge, and size of the ion.
Cations are smaller than anions. The charge on cations is more concentrated. Due to this, the interaction between a cation and negative end of the polar molecule are stronger than the corresponding interaction between the anion and positive end of the polar molecule.
When NaCl is dissolved in water it dissociates into Na+ and Cl– ions. They undergo hydration.
Hydration is a process by which each ion is surrounded by a number of water molecules with an oppositely charged pole of solvent water oriented in the direction of the ion concerned.
Dispersion Forces or London Forces due to Interaction Between Molecules:
Atoms and nonpolar molecules are electrically symmetrical and have no dipole moment because their electronic charge cloud is symmetrically distributed. But a dipole may develop momentarily even in such atoms and molecules due to polarizability. Polarizability is an ease with which the arrangement of electrons in the atom or molecule can be disturbed.
Suppose we have two nonpolar atoms ‘A’ and ‘B’ in the close vicinity of each other. It may so happen that momentarily electronic charge distribution in one of the atoms say ‘A’, becomes unsymmetrical i.e., the charge cloud is more on one side than the other.
This results in the development of instantaneous dipole on the atom ‘A’ for a very short time. This instantaneous or transient dipole distorts the electron density of the other atom ‘B’, which is close to it and as a consequence, a dipole is induced in the atom ‘B’. The temporary dipoles of atom ‘A’ and ‘B’ attract each other. Similarly temporary dipoles are induced in molecules also.
This force of attraction was first proposed by the German physicist Fritz London, and for this reason force of attraction between two temporary dipoles is known as London forces. Another name for this force is dispersion force.
These forces are always attractive and interaction energy is inversely proportional to the sixth power of the distance between two interacting particles. where r is the distance between two particles.
These forces are important only at short distances (~500 pm) and their magnitude depends on the polarizability of the particle.
The strength of these forces increases with the increase in molecular mass, molecular size, number of electrons and surface area of the molecule.
The liquid state of helium and methane is due to the presence of dispersion forces. Other examples are N2, H2, CO2.
Hydrogen Bonding Between Molecules:
Consider two polar molecules water H2O (molecular mass 18 and dipole moment 1.8 D) and nitrosyl fluoride ONF (molecular mass 49 and dipole moment 1.8 D).
As their dipole moments are the same, their boiling points should be comparable. But boiling point of water is 100° C while that of nitrosyl fluoride is – 56° C. Thus there is a large difference between their boiling points. This difference is due to a special type of bonding between water molecules called hydrogen bonding.
The large difference in electronegativities of the pairs of bonded atoms N-H, O-H, F-H and Cl-H establishes highly polar covalent bonds in which hydrogen acquires exceptionally large positive charge and another electronegative atom acquires exceptionally large negative charge on it.
Due to this, there is an attraction between the positively charged hydrogen atom of one molecule and negatively charged atom of another molecule. This interaction is called hydrogen bonding.
This interaction is represented by a dotted line. The energy of hydrogen bond varies between 10 to 100 kJ mol-1. This energy is significant. Hence it is a very important factor which influences bulk properties of a substance.
Variation of Boiling Points of Hydrides:
Following graph shows variation of boiling points of hydrides of groups 4A, 5A, 6A, 7A.
From graph, we can see that the boiling point of ammonia (hydride of nitrogen), water (hydride of oxygen) and hydrogen fluoride (hydride of fluorine) are exceptionally higher in their group. It is due to strong molecular hydrogen bonding.
We can see that methane (hydride of carbon) has low boiling point due to the absence of hydrogen bonding.
In case of inert gases, the trend is increasing in boiling points with the increase in atomic mass. It is due to dispersion forces increase with the mass of the molecule.