Sublimation (phase transition)

Sublimation is the transition of a substance directly from the solid phase to the gas phase without passing through an intermediate liquid phase. 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., water ice just below 0C). 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. Snow and ice sublime, although more slowly, below the melting point temperature.[1] This allows a wet cloth to be hung outdoors in freezing weather and retrieved later in a dry state. In freeze-drying, the material to be dehydrated is frozen and its water is allowed to sublime under reduced pressure or vacuum. The loss of snow from a snowfield during a cold spell is often caused by sunshine acting directly on the upper layers of the snow. Ablation is a process that includes sublimation and erosive wear of glacier ice. The melting point of a solid is the temperature at which it changes state from solid to liquid. At the melting point the solid and liquid phase exist in equili

rium. The melting point of a substance depends (usually slightly) on pressure and is usually specified at standard pressure. When considered as the temperature of the reverse change from liquid to solid, it is referred to as the freezing point or crystallization point. Because of the ability of some substances to supercool, the freezing point is not considered as a characteristic property of a substance. When the "characteristic freezing point" of a substance is determined, in fact the actual methodology is almost always "the principle of observing the disappearance rather than the formation of ice", that is, the melting point.[1] For most substances, melting and freezing points are approximately equal. For example, the melting point and freezing point of the element mercury is 234.32 kelvin (?38.83 C or ?37.89 F). However, certain substances possess differing solid-liquid transition temperatures. For example, agar melts at 85 C (185 F) and solidifies from 31 C to 40 C (89.6 F to 104 F); such direction dependence is known as hysteresis. The melting point of ice at 1 atmosphere of pressure is very close [2] to 0 C (32 F, 273.15 K); this is also known as the ice point. In the presence of nucleating substances the freezing point of water is the same as the melting point, but in the absence of nucleators water can supercool to ?42 C (?43.6 F, 231 K) before freezing. The chemical element with the highest melting point is tungsten, at 3683 K (3410 C, 6170 F) making it excellent for use as filaments in light bulbs. The often-cited carbon does not melt at ambient pressure but sublimes at about 4000 K; a liquid phase only exists above pressures of 10 MPa and estimated 43004700 K. Tantalum hafnium carbide (Ta4HfC5) is a refractory compound with a very high melting point of 4488 K (4215 C, 7619 F).[3] At the other end of the scale, helium does not freeze at all at normal pressure, even at temperatures very close to absolute zero; pressures over 20 times normal atmospheric pressure are necessary.