Intermolecular Forces

States Of Matter (Intermolecular Forces)

We know that matter exists in different states: solid, liquid, and gas. These different states arise because of the interplay between two opposing forces or energies:

The attractive forces between the particles (atoms, molecules, or ions). These are called intermolecular forces (or interparticle forces).
The energy associated with the motion of particles, called thermal energy, which tends to keep the particles apart.

The state of matter is a result of the balance between these two factors. Stronger intermolecular forces and lower thermal energy favour the condensed states (solid and liquid), where particles are held close together. Weaker intermolecular forces and higher thermal energy favour the gaseous state, where particles are far apart and move randomly.

Intermolecular Forces

Intermolecular forces are the attractive or repulsive forces that act between neighbouring particles (atoms, molecules, or ions). It is important to distinguish these from intramolecular forces, which are the strong forces (covalent bonds, ionic bonds, metallic bonds) holding atoms together *within* a molecule or compound.

Intermolecular forces are generally much weaker than intramolecular forces (covalent or ionic bonds). However, they are crucial in determining the physical properties of substances, such as melting point, boiling point, solubility, viscosity, and surface tension.

Different types of intermolecular forces exist, depending on the nature of the particles involved. They are collectively known as van der Waals forces (except for Hydrogen bonding, which is considered a special, stronger type).

Dispersion Forces Or London Forces

These forces are present between all types of particles, including non-polar molecules and even individual atoms (like noble gases). They are the weakest type of intermolecular force.

Origin: Dispersion forces arise from the instantaneous, temporary fluctuations in electron distribution within an atom or molecule. Even in a non-polar molecule with symmetrically distributed electrons, at any given instant, there might be a momentary imbalance, creating a temporary or instantaneous dipole.

Illustration of temporary dipole formation

This instantaneous dipole in one particle can induce a temporary dipole in a neighbouring particle by distorting its electron cloud. A weak attractive force then exists between these two temporary dipoles.

Illustration of induced dipole and dispersion forces

These dipoles are constantly forming and disappearing, but their instantaneous interactions lead to a net weak attraction.

Strength: The strength of dispersion forces depends on the polarisability of the particles (how easily the electron cloud can be distorted) and the contact area between molecules. Larger molecules with more electrons and greater surface area are more polarisable and experience stronger dispersion forces. This explains why, for example, larger noble gases (like Xenon) have higher boiling points than smaller noble gases (like Helium), or why straight-chain hydrocarbons have higher boiling points than branched isomers of the same molecular weight.

Dipole - Dipole Forces

These forces exist between molecules that have permanent dipoles. A permanent dipole occurs in polar molecules due to the unequal sharing of electrons in polar covalent bonds, resulting in a separation of positive and negative charges.

Origin: The positive end of one polar molecule is attracted to the negative end of a neighbouring polar molecule. This electrostatic attraction constitutes the dipole-dipole force.

Illustration of dipole-dipole forces between polar molecules

Strength: Dipole-dipole forces are generally stronger than dispersion forces for molecules of comparable size and mass, because they involve permanent, rather than temporary, dipoles. The strength of dipole-dipole forces increases with increasing polarity of the molecules (i.e., larger dipole moment).

Dipole-dipole forces are additive to dispersion forces; polar molecules experience both types of attractions.

Example: HCl (Hydrogen chloride). The chlorine atom is more electronegative than hydrogen, creating a partial negative charge ($\delta^-$) on Cl and a partial positive charge ($\delta^+$) on H. Molecules align such that the $\delta^+$ of one is near the $\delta^-$ of another.

Dipole–Induced Dipole Forces

These forces occur between a polar molecule (with a permanent dipole) and a non-polar molecule (which can have a temporary dipole induced in it).

Origin: The permanent dipole of the polar molecule can distort the electron cloud of the neighbouring non-polar molecule, inducing a temporary dipole in it. An attractive force then exists between the permanent dipole and the induced dipole.

Illustration of dipole-induced dipole forces

Strength: The strength of dipole-induced dipole forces depends on the magnitude of the permanent dipole of the polar molecule and the polarisability of the non-polar molecule. These forces are generally weaker than dipole-dipole forces but stronger than pure dispersion forces for molecules of similar size.

Example: When a polar molecule like water is mixed with a non-polar molecule like oxygen, the water molecule's dipole induces a dipole in the oxygen molecule, allowing a weak attraction to form. This force is responsible for the solubility of non-polar gases like oxygen and nitrogen in water.

Hydrogen Bond

Hydrogen bond is a special, strong type of dipole-dipole interaction. It occurs when a hydrogen atom is bonded to a highly electronegative atom (such as Nitrogen (N), Oxygen (O), or Fluorine (F)) and is simultaneously attracted to a lone pair of electrons on another electronegative atom (N, O, or F) in a neighbouring molecule (or sometimes within the same large molecule).

Conditions for Hydrogen bonding:

Presence of a highly polar H-X bond, where X is a small, highly electronegative atom (N, O, or F). This makes the Hydrogen atom significantly positive ($\delta^+$).
Presence of another electronegative atom (N, O, or F) with at least one lone pair of electrons, which can act as a hydrogen bond acceptor.

Illustration of Hydrogen bonding in water molecules

Example: Water (H$_2$O). Oxygen is highly electronegative, creating a significant $\delta^+$ on the two hydrogen atoms and $\delta^-$ on the oxygen atom. The $\delta^+$ hydrogen of one water molecule is attracted to the lone pair on the oxygen of a neighbouring water molecule.

Strength: Hydrogen bonds are significantly stronger than typical dipole-dipole and dispersion forces (though still much weaker than covalent bonds). A typical hydrogen bond energy is around 10-40 kJ/mol, compared to a few kJ/mol for dispersion forces and dipole-dipole forces, and 200-800 kJ/mol for covalent bonds.

Significance: Hydrogen bonding has profound effects on the properties of substances:

Unusually high boiling points of water, ammonia, and hydrogen fluoride compared to heavier analogues (like H$_2$S, PH$_3$, HCl).
High specific heat and heat of vaporisation of water.
Lower density of ice compared to liquid water (due to the open cage-like structure formed by hydrogen bonds in ice).
Structure of biological molecules like DNA (holding the two strands together) and proteins (in secondary structures like alpha-helix and beta-pleated sheet).

Thermal Energy

Thermal energy is the energy possessed by the particles of matter due to their motion. It is a measure of the average kinetic energy of the particles.

Thermal energy is directly proportional to temperature. Higher the temperature, the more vigorously the particles move, and the greater is their thermal energy.

The motion of particles can be:

Translational motion: Movement of particles from one place to another (dominant in gases and liquids).
Vibrational motion: Oscillations of particles about their mean positions (present in all states, dominant in solids).
Rotational motion: Rotation of particles about their axis (present in gases and liquids).

Thermal energy tends to keep the particles apart. It causes randomness and disorder in the system.

Intermolecular Forces Vs Thermal Interactions

The physical state of a substance is a result of the competition between the attractive intermolecular forces and the disruptive thermal energy.

Intermolecular forces: Tend to bring the particles closer together and hold them in fixed or relatively fixed positions, leading to order (solid state) or close proximity (liquid state).
Thermal energy: Tends to move the particles apart and create randomness and disorder (gaseous state).

The relative strength of these two factors determines the state of matter at given conditions of temperature and pressure:

Solids: At low temperatures, thermal energy is low, and intermolecular forces are strong enough to hold particles in fixed positions, resulting in a rigid structure. Intermolecular forces are much stronger than thermal energy.
Liquids: As temperature increases, thermal energy increases. Particles gain enough energy to overcome the strong intermolecular forces holding them in fixed positions, but the forces are still strong enough to keep them close together. Thermal energy and intermolecular forces are comparable in magnitude. Particles can move past each other, resulting in fluidity but maintaining a definite volume.
Gases: At high temperatures, thermal energy is much greater than the intermolecular forces. Particles move randomly at high speeds, far apart from each other. Intermolecular forces are negligible compared to thermal energy. The gas expands to fill the entire volume.

By changing the temperature and pressure, we can alter the balance between thermal energy and intermolecular forces, causing matter to undergo phase transitions (e.g., melting, boiling, condensation, freezing, sublimation).

For instance, heating a solid increases the thermal energy of its particles. When the thermal energy becomes sufficient to overcome the forces holding the particles in fixed positions, the solid melts to form a liquid. Further heating increases thermal energy, and when it becomes much larger than the intermolecular forces, the liquid boils to form a gas.