Sublimation is the process of transition of a substance from the solid phase to the gas phase without passing through an intermediate liquid phase (can also occur in opposite order, such as in Hoar frost). Sublimation is an endothermic phase transition that occurs at temperatures and pressures below a substance's triple point in its phase diagram.
At normal pressures, most chemical compounds and elements possess three different states at different temperatures. In these cases, the transition from the solid to the gaseous state requires an intermediate liquid state. Note, however, that the pressure referred to here is the partial pressure of the substance, not the total (e.g., atmospheric) pressure of the entire system. So, all solids that possess an appreciable vapor pressure at a certain temperature usually can sublime in air (e.g., ice just below 0°C). For some substances, such as carbon and arsenic, sublimation is much easier than evaporation from the melt, because the pressure of their triple point is very high, and it is difficult to obtain them as liquids.
Sublimation requires additional energy and is an endothermic change. The enthalpy of sublimation (also called heat of sublimation) can be calculated as the enthalpy of fusion plus the enthalpy of vaporization. The reverse process of sublimation is deposition. The formation of frost is an example of meteorological deposition.
Condensation is initiated by the formation of atomic/molecular clusters of that species within its gaseous volume—like rain drop or snow-flake formation within clouds—or at the contact between such gaseous phase and a (solvent) liquid or solid surface.
A few distinct reversibility scenarios emerge here with respect to the nature of the surface.
- absorption into the surface of a liquid (either of the same species or one of its solvents)—is reversible as evaporation.