Intermolecular forces are the forces of attraction or repulsion that may exist between molecules that are in close vicinity to each other. These forces are responsible for physical properties like boiling point, melting point, density, vapor pressure, viscosity, surface tension, and solubility of compounds. These forces are much weaker than intramolecular/interatomic forces. Interatomic or intramolecular forces act between atoms and result in the formation of chemical bonds. Intermolecular forces can be categorized into two main types:

• Dipole-dipole interactions
• Van der Waals forces

These interactions are formed due to uneven distribution of electrons in a molecule. This gives rise to a partial positive (+δ) and a partial negative (-δ) charge in a molecule that, as a whole, is neutral. Polar molecules tend to orient themselves in such a way that the +δ part of the molecule is close to the -δ part of the molecule, such that there is minimum repulsion and maximum attraction between the molecules. Dipole-dipole interactions can be further categorized into three types:

1. Ion-dipole forces
2. Ion-induced dipole forces
3. Hydrogen bonds

The force of attraction between a polar molecule and an ion that may lie in its vicinity is called an ion-dipole force. Example

• When NaCl is dissolved in water, it will dissociate into Na⁺ ions and Cl^- ions; the force of attraction that may exist between, say, Na⁺ and the -δ oxygen of water is nothing but ion-dipole force. It is due to this force of attraction that the polar molecule will dissolve in a polar solvent like water.

The force of attraction between a non-polar molecule and an ion that may lie in its vicinity is called ion-induced dipole force. In this, the ion may attract or repel the electron cloud present on the non-polar molecule and induce the non-polar molecule to become a temporary dipole. The strength of this induced dipole depends on how easily the electron cloud can be distorted, i.e., the bigger the molecule, the stronger is the dipole induced. Example

• Hemoglobin is a protein found in the red blood cells, and its function is to carry oxygenated blood to various parts of the body. It has an Fe²⁺ ion in the center of its protein structure. This Fe²⁺ ion attracts the O₂ by ion-induced dipole force. (Although oxygen is an electronegative atom, in O₂, the electron pairs experience an equal pull from both the oxygen atoms, and thus, there is no development of +δ and -δ charge on O₂. Hence, the molecule as a whole is non-polar).

The force of attraction between the lone pair of electrons in an electronegative atom (atoms in a covalent bond that tend to pull the shared pair of electrons towards themselves) and a hydrogen atom that is covalently attached to either nitrogen, fluorine, or oxygen is called a hydrogen bond. The hydrogen atom is attached to either nitrogen, fluorine or oxygen, and all these atoms are more electronegative than hydrogen. They tend to pull the shared pair of electrons towards themselves and develop a -δ charge. The hydrogen atom, on the other hand, develops a +δ charge on itself. Now, the molecules tend to orient themselves in such a way that the +δ hydrogen atom is close to the electronegative atom, and the force of attraction that develops between the lone pair of electrons (in the electronegative atom) and the +δ hydrogen atom is called a hydrogen bond. Examples

• In water, there exists a hydrogen bond between the electronegative oxygen of one water molecule and the +δ hydrogen atom of another water molecule. These are the most prominent intermolecular forces acting in water.
• In ammonia, there exists a hydrogen bond between the lone pair electrons of nitrogen of one ammonia molecule and the +δ hydrogen atom of another ammonia molecule. Here, nitrogen has only one lone pair of electrons, whereas in oxygen, there are two lone pairs of electrons; therefore, the strength of hydrogen bond in water is much greater than that compared to ammonia.

Van der Waals forces are usually the forces of attraction and repulsion that may exist between molecules and surfaces. They are weaker than a hydrogen bond. These forces are dependent on the orientation of the molecule. Usually, they are weak forces of attraction that exist between neutral molecules. These forces can act on longer distances as compared to other intermolecular forces of attraction. However, these forces do not act beyond a particular distance. As the molecules come closer, the van der Waals forces of attraction keep on increasing until they reach a particular level of proximity called van der Waals contact distance. Beyond this distance, van der Waals forces of attraction keep on decreasing as the forces of repulsion between the molecule increases. This is because their outer electron clouds overlap. It is due to van der Waals forces that real gases deviate from their ideal gas properties; this deviation from the ideal gas properties can be explained by van der Waals equation given below, which takes into account the volume occupied by the molecules of gas and also the force of attraction that may exist between them, i.e., the van der Waals forces. where, n = Number of moles of gas p = Pressure exerted by the gas T = Absolute temperature of the system V = Total volume of the gas in the container R = Universal gas constant a = Na² = a'= Total force of attraction that exists between all the particles in mole one of the gas b = Na.b' = Total volume occupied by one mole of particles of the gas (Note: For an ideal gas, the above equation can be written as PV = nRT) Van der Waals forces are prominent in molecules where other intermolecular forces do not exist. They can further be classified into three other types:

1. Keesom interactions
2. Debye force
3. London dispersion force

These interactions occur between permanent dipoles, which can be either molecular ions, dipoles (polar molecules) or quadrupoles (e.g. in CCl₄, the electrons of the carbon atom experience an equal pull in all four directions, and hence, the molecule as a whole is non-polar). These molecules tend to orient themselves in such a way that the +δ part of the molecule is close to the -δ part of the molecule; thus, there is minimum repulsion and maximum attraction between the molecules. These interactions are temperature-dependent. They tend to account for both forces of attraction and repulsion that may exist between two molecules. Example

• HCl is a polar molecule. When two HCl molecules come closer, they tend to orient themselves in such a way that there is maximum force of attraction and minimum repulsion between them.

These interactions occur between permanent dipoles and induced dipoles. In other words, it is the interactions that occur between a polar molecule and a molecule that can be polarized in the presence of a polar molecule. Ease of polarization of molecules increases with the size of the electron cloud and thus, the size of the molecule. The polar molecule tends to shift (usually repel) the non-polar molecule's electron cloud to one side of the molecule, giving rise to an induced polarity. Again, the molecules tend to orient themselves in such a way that there is maximum force of attraction between the molecules. Debye force usually accounts for only the forces attraction acting between molecules. Example

• Argon and HCl: HCl is a polar molecule. When the non-polar argon atom and HCl come closer, the -δ part of HCl repels the electron cloud, which then shifts to side of the atom and induces argon to become temporarily polar. Both the molecules orient themselves in such a way that there is maximum attraction and minimum repulsion between the molecule.

This kind of force arises due to the instantaneous dipole that may be created in the atoms of molecules due to the movement of electrons. As the electrons in an atoms are in continuous motion, there might be an instance when most of the electrons have shifted to one side of the electron cloud causing a momentary dipole to be created. When two such instantaneous dipoles come close together, there is attraction between the molecules. This is nothing but London dispersive force. This is the weakest amongst all the forces, but is present in almost all molecules and atoms. Example

• Ne and Ne: When two momentary dipoles of neon come close, there is a force of attraction that acts between them.

In nature, there may be one or more than one intermolecular forces that may act on a molecule. Let us look at the following examples to get a better understanding of it. Intermolecular Forces Acting on Water Water is a polar molecule, with two +δ hydrogen atoms that are covalently attached to a -δ oxygen atom. Thus, the water molecule exhibits two types of intermolecular forces of attraction. These are hydrogen bonds and London dispersion force. Intermolecular Forces in NH₃ In NH₃, there is a -δ nitrogen that is covalently attached to three +δ hydrogen atoms. Thus, the ammonia molecule exhibits two types of intermolecular force of attraction. These are hydrogen bonds and London dispersion force. Intermolecular Forces in CH₄ CH₄ is a symmetric non-polar molecule, and thus, it exhibits only London dispersion force.