What is the vapor pressure of a substance? How does it depend on the temperature and strength of intermolecular forces?

In the context of vapor pressure, thermodynamic equilibrium refers to a balanced state where the rate of evaporation of a liquid (or sublimation of a solid) equals the rate of condensation from its vapor. This balance occurs at a specific pressure for a given temperature in a closed system. When a substance reaches this equilibrium, the air space above the substance in a sealed container becomes saturated with vapor particles, and no net change is observed between the vapor and its condensed phase.
In simpler terms, imagine a party where some guests are inside and some outside. As guests enter and leave the house at the same rate, the number of people inside stays the same, just like the amount of vapor and liquid remain constant at equilibrium.

The strength of intermolecular forces — attractions between molecules — significantly impacts a substance's vapor pressure. These forces include hydrogen bonds, dipole-dipole interactions, and Van der Waals forces, which vary in strength. For instance, hydrogen bonds are much stronger than Van der Waals forces.

To visualize this, consider a group of people holding hands (molecules in a liquid). If they hold hands tightly (strong intermolecular forces), it's difficult for an individual to leave the group (low vapor pressure). Contrastly, if they barely touch fingertips (weak intermolecular forces), it's much easier to let go and move away (high vapor pressure). This concept is vital in understanding why substances like water, which forms strong hydrogen bonds, have lower vapor pressures compared to those that only exhibit weak Van der Waals forces.

The relationship between vapor pressure and temperature is direct and crucial. As temperature increases, so does the vapor pressure of a substance. This is because temperature is a measure of the average kinetic energy of the molecules. At higher temperatures, more molecules gain sufficient energy to break free from the intermolecular forces that hold them in the liquid or solid phase.

Consider our party analogy again: as the evening progresses (temperature increases), the energy and excitement among guests rise (kinetic energy). More guests may decide to leave the inside of the house to enjoy the outside (vapor phase), suggesting an increase in the 'vapor pressure' of the house party atmosphere. An understanding of this principle is essential for explaining phenomena such as boiling, where the vapor pressure equals atmospheric pressure, allowing bubbles of vapor to form within the liquid.