Intermolecular Forces

Liquids have qualities that are halfway between gases and solids, yet they resemble solids more. Intermolecular forces hold molecules together in a liquid or solid, as opposed to intramolecular forces like covalent bonds that hold atoms together in molecules and polyatomic ions. Covalent bonds are stronger than intermolecular forces. In 1 mol of water, for example, it takes 927 kJ to overcome intramolecular forces and break both O–H bonds, while it only takes roughly 41 kJ to overcome intermolecular attractions and convert 1 mol of liquid water to water vapour at 100°C (Despite this low value, liquid water’s intermolecular forces are among the strongest known). Changes between the solid, liquid, and gaseous states almost often occur for molecular compounds due to the huge variation in the strengths of intramolecular and intermolecular forces.

Liquids have qualities that are halfway between gases and solids, yet they resemble solids more.

The melting and boiling points of solids and liquids are determined by intermolecular forces. When the molecules in a liquid have enough thermal energy to overcome the intermolecular attraction interactions that keep them together, bubbles of vapour develop. Solids melt when molecules gain enough thermal energy to overcome the intermolecular interactions that keep them locked in place.

Intermolecular forces are electrostatic in origin, meaning they are caused by the interaction of positively and negatively charged species. Intermolecular interactions are made up of both attracting and repulsive components, just like covalent and ionic connections. Intermolecular interactions are most essential for solids and liquids, where the molecules are close together, because electrostatic interactions decrease rapidly with increasing distance between molecules. Only at very high pressures do these interactions become significant for gases, and they are responsible for observable deviations from the ideal gas law.

Three types of intermolecular interactions are explicitly discussed in this section. Ion–ion interactions, which are important for ionic bonding, and ion–dipole interactions, which occur when ionic substances dissolve in a polar substance like water, are two other types of electrostatic interactions that you are already familiar with. Van der Waals forces are a term used to describe the first two.

Dipole Dipole Interaction

The linked atoms behave as if they have equal but opposite localised fractional charges (i.e., the two bonded atoms generate a dipole). A molecule possesses a net dipole moment if its structure prevents the individual bond dipoles from cancelling one another. Molecules with net dipole moments prefer to arrange themselves in such a way that the positive end of one dipole is near the negative end of another, as seen:

Dipole–Dipole Interactions: Attractive and Repulsive. (a and b) Attractive interactions are produced when the positive end of one dipole (+) is near the negative end of another (and vice versa). (c and d) Repulsive interactions are produced when the positive or negative ends of dipoles on neighbouring molecules are juxtaposed

These configurations are more stable than those with two neighbouring positive or negative terminals. As a result, dipole–dipole interactions like those in (b) are attractive intermolecular interactions, whereas those shown in (d) are repulsive. Because molecules in a liquid move freely and continuously, they are constantly subjected to both attractive and repulsive dipole–dipole interactions. The attractive interactions, on the other hand, tend to win out in the end.

Dipole–dipole interactions are far weaker than interactions between two ions, each of which has a charge of at least 1, or between a dipole and an ion, each of which has at least a complete positive or negative charge. Furthermore, the attractive attraction between dipoles decays significantly faster than ion–ion interactions as the distance between them increases. Remember that the attractive energy between two ions is proportional to 1/r, where r is their distance. The attracting energy is cut in half by doubling the distance (r=2r). The energy of a dipole-dipole interaction, on the other hand, is proportional to 1/r3, so doubling the distance between the dipoles reduces the strength of the contact by 2,3, or eight times. At ambient temperature and 1 atm pressure, a substance like HCl is a gas because it is partially held together by dipole–dipole interactions. NaCl, on the other hand, is a solid with a high melting point because of interionic interactions. The intensity of intermolecular interactions grows as the dipole moment of the molecules increases within a succession of compounds with identical molar masses.

Conclusion 

We conclude that Many of a substance’s qualities are determined by intermolecular forces, which hold numerous molecules together. Van der Waals forces refer to all attractive forces between neutral atoms and molecules, albeit they are more commonly referred to as intermolecular attraction.