Sublimation is the direct phase transition in which a substance changes from a solid to a gas without passing through the liquid state under the conditions being considered. It is a physical change because the chemical identity of the substance is ordinarily preserved. Familiar examples include dry ice, solid carbon dioxide, changing directly into carbon dioxide gas, and certain forms of ice gradually disappearing into water vapour under sufficiently dry conditions.
The reverse process, in which a gas changes directly into a solid, is called deposition. Sublimation is favoured when the pressure and temperature conditions make the gas phase thermodynamically accessible from the solid while the liquid phase is not stable or is bypassed under the particular conditions. The term comes from historical descriptions of substances apparently rising directly into vapour, but modern thermodynamics defines it precisely as a solid to gas phase transition. Sublimation should not be confused with evaporation, because evaporation is a liquid to gas process occurring at a liquid surface, whereas sublimation begins with a solid.
At the microscopic level, sublimation occurs when particles at or near the surface of a solid acquire enough energy and an appropriate momentum distribution to escape the cohesive forces holding them within the solid. Even below the conventional sublimation temperature, a solid has particles with a range of energies, so some surface molecules can escape into the gas phase. As molecules leave, the solid loses particles directly to the vapour.
The rate depends on temperature, vapour pressure, surface area, ambient pressure, surrounding gas composition, molecular interactions and the material's crystal structure. Increasing temperature generally increases the equilibrium vapour pressure of a solid and therefore increases sublimation tendency. Lowering the surrounding pressure can also make direct solid to gas transition easier because fewer gas molecules oppose the escaping particles and because the thermodynamic conditions may move into the solid gas region of the phase diagram. Sublimation is consequently both a molecular transport process and a thermodynamic phase transition.
Phase diagrams are essential for understanding when sublimation can occur. A substance has characteristic boundaries separating its solid, liquid and gas regions, and the solid gas boundary represents equilibrium between the solid and gaseous phases. The triple point is particularly important because below the pressure of the triple point, a stable liquid phase cannot exist in equilibrium. Under such conditions, heating the solid can cause it to sublime rather than melt and subsequently boil.
Carbon dioxide provides the classic practical example: at ordinary atmospheric pressure, solid carbon dioxide does not form a stable liquid phase and instead changes directly to gas at approximately minus 78.5 degrees Celsius. Water can also sublime, especially under low vapour pressure conditions. In a sufficiently cold and dry environment, snow or ice can disappear without producing a visible liquid layer because water molecules enter the atmosphere directly as vapour. This contributes to snow loss in some polar and mountainous environments even when temperatures remain below freezing.
Sublimation requires energy because particles must move from a condensed solid state into the much more dispersed gas phase. The relevant enthalpy is called the enthalpy of sublimation, and for a substance undergoing the transition at equilibrium it represents the energy required per mole or per unit mass for solid to gas conversion under specified conditions.
Thermodynamically, the enthalpy of sublimation is related to the enthalpies of fusion and vaporisation through Hess's law when the relevant phases and conditions are appropriately defined. Because sublimation involves a much larger increase in molecular freedom than melting, its entropy change is generally substantial. Heat supplied to a sublimating solid can therefore be consumed by the phase transition rather than simply increasing the temperature. If the surroundings supply insufficient energy, sublimation slows. If a sublimating system is isolated, energy consumption by escaping molecules can cool the remaining solid or its surroundings. This cooling effect contributes to the familiar behaviour of dry ice and some other volatile solids.
Sublimation has important industrial and scientific applications. Freeze drying, or lyophilisation, removes water from frozen material by reducing pressure so that ice sublimes rather than melting and then evaporating. This process is widely used for preserving pharmaceuticals, biological materials, foods and other temperature sensitive products because removing water at low temperatures can help preserve structure and activity. Sublimation is also used in purification and separation techniques because different substances have different vapour pressures and sublimation behaviours.
In chemistry, sublimation can sometimes separate a volatile solid from nonvolatile impurities. In semiconductor and materials processing, vacuum deposition methods exploit closely related vapour transport principles to move material from a condensed source to another surface. Solid carbon dioxide is used for cooling, cleaning and controlled atmospheric effects because it sublimes under ordinary pressure without leaving a liquid residue. The useful properties of sublimation therefore arise directly from the combination of phase equilibrium, volatility and controllable pressure.
Sublimation is influenced strongly by vapour pressure and ambient conditions. A solid has an equilibrium vapour pressure that generally rises rapidly with temperature, and sublimation occurs when molecules escape faster than they are returned by condensation or when the environment is undersaturated with respect to the solid. In a closed system containing a solid and its vapour, an equilibrium can eventually be established in which sublimation and deposition occur at equal rates.
In an open environment, continuous removal of vapour can maintain net sublimation. Air movement, humidity and atmospheric pressure therefore matter. Ice sublimates more readily in dry air than in humid air because dry air has a greater capacity to accept additional water vapour. Wind can enhance mass transfer by removing vapour near the surface. Surface area also matters because sublimation is a surface process, so powdered or porous materials can sublime more rapidly than compact pieces under otherwise comparable conditions.
Sublimation has several important connections with astronomy and planetary science. On bodies with very low atmospheric pressures, volatile substances can transition directly between solid and gas phases. Water ice, carbon dioxide ice, nitrogen ice and other frozen volatiles can participate in sublimation and deposition cycles on planets, moons, comets and asteroids.
Cometary activity is strongly influenced by solar heating of volatile ices beneath or at the surface, producing gases that can carry dust away and contribute to the visible coma and tail. On Mars, seasonal carbon dioxide frost can deposit from the atmosphere and later sublime as sunlight returns. Similar processes affect the surfaces and atmospheres of other cold planetary bodies. These examples show that sublimation is not an unusual laboratory curiosity but a fundamental mechanism governing the movement of volatile material where temperature and pressure favour direct solid gas exchange.
A common misconception is that sublimation means any solid gradually disappearing. A solid may also be consumed by chemical decomposition, dissolution, abrasion or reaction with the surrounding environment, none of which is necessarily sublimation. Another misconception is that sublimation requires extremely high temperatures; in reality, sublimation can occur at low temperatures whenever the solid has a nonzero vapour pressure and the surrounding conditions permit net molecular escape.
Sublimation can coexist with other processes, including evaporation, melting, freezing and deposition, depending on local temperature and pressure. In thermodynamic terms, the solid to gas transition becomes favourable when the gas phase has lower Gibbs free energy than the solid under the relevant conditions. The increase in entropy is an important component of this transition because gas particles have vastly more accessible spatial configurations than particles constrained within a solid. Understanding sublimation therefore requires considering molecular energy distributions, surface transport, vapour pressure, phase diagrams and the thermodynamic balance between phases rather than treating it simply as an unusual form of evaporation.
