Sublimation is a process of transformation directly from the solid phase into the gaseous phase whithout going through an intermediate phase such as liquid phase. This endohermic phase transition occurs at temperatures and pressures below a substance’s triple point in its phase diagram.
When the pressure is optimal, most of the chemical compounds and elements possess three different states at different temperatures. In these cases, when the transition takes place, it is required an intermediate liquid state. It’s important to know that the pressure mentioned here is the partial pressure of the substance, not the total (e.g.: atmospheric) pressure of the entire system. Therefore, all solids that possess an appreciable vapor pressure at a certain temperature usually can sublime in air (e.g.: water ice just below 0°C). If talking about substances like carbon or arsenical, 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.
The sublimation process needs 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.
Solid carbon dioxide (dry ice) sublimes readily at atmospheric pressure at -78.5°C (197.5 K, −104.2 °F), whereas liquid CO2 can be obtained at pressures and temperatures above the triple point (5.2 atm, -56.4°C).