Phase Transitions: Sublimation and Deposition

Sublimation is the direct transition of a solid to gas without passing through the liquid phase. Unlike melting, which converts solid to liquid, sublimation completely bypasses the liquid state. Dry ice exemplifies this process, transforming directly into carbon dioxide vapor at −78.5 °C. Sublimation occurs in solids with weak intermolecular forces that cannot maintain the solid structure when sufficient thermal energy is applied.

Solids exhibiting weak intermolecular forces require less energy to overcome attractive forces between molecules. In dry ice, weak dispersion forces between CO2 molecules allow surface molecules to escape directly into the vapor phase when they acquire sufficient thermal energy. Stronger intermolecular forces would require more energy, making the solid more stable and preventing sublimation under normal conditions.

The molar heat of sublimation, or molar enthalpy of sublimation, is the energy required to convert one mole of solid directly into gas. This value is always positive because sublimation is an endothermic process requiring energy input to overcome intermolecular attractions. The energy must be supplied to break bonds holding molecules in the solid state and allow them to escape as gas.

Deposition is the reverse of sublimation—the direct transition of gas to solid without forming liquid. When gas molecules collide with cooler solid surfaces, they lose heat and deposit onto the surface. Frost formation exemplifies deposition. While sublimation is endothermic with positive enthalpy, deposition is exothermic with negative enthalpy of equal magnitude.

In open systems, sublimed molecules disperse into air and rarely return, making sublimation rate exceed deposition rate. In closed systems, a solid-vapor equilibrium establishes at the sublimation point, where sublimation and deposition rates balance. This equilibrium creates a dynamic state where gas molecules continuously deposit back onto the solid surface.

Vapor pressure is the partial pressure exerted by gas in dynamic equilibrium with its solid at the sublimation point. Solids that sublimate have high vapor pressures because their weak intermolecular forces allow molecules to escape easily into the gas phase. Dry ice, for example, exhibits a vapor pressure of 56.5 atm at 20 °C, far exceeding most ordinary solids.

Sublimation can be modeled as sequential melting followed by vaporization using Hess's Law. The enthalpy of sublimation approximately equals the sum of the enthalpy of fusion and enthalpy of vaporization. This approach works because converting solid to gas requires completely overcoming intermolecular attractions, whereas melting only partially overcomes them.