Superheating is persistence of a phase above the equilibrium temperature at which a higher-temperature phase becomes favorable. The familiar case is a liquid above its boiling point at the imposed pressure, although delayed melting can also be discussed as superheating.
The liquid can remain in a metastable state because producing a vapor bubble costs interfacial energy. A sufficiently large bubble overcomes the nucleation barrier and grows, converting the liquid and consuming latent heat. Smooth clean containers can delay nucleation, while cavities and boiling chips provide heterogeneous nucleation sites. The observed temperature of sudden bubble formation therefore need not equal the equilibrium boiling point. The amount of superheating is a kinetic and preparation-dependent observation, rather than a new equilibrium phase boundary.
A liquid maintained above its equilibrium boiling temperature at the imposed pressure. A barrier to nucleation can delay its conversion into vapor, making this a metastable state.
A solid can locally exceed its solidus because a thin solute-diffusion layer changes composition faster than thermal conduction changes temperature. Near a melting phase boundary, a depleted solid can join the equilibrium solidus at the interface while the adjacent, more concentrated solid remains above its own lower solidus. The assumed single sharp interface is then thermodynamically unstable to additional melting and may be replaced by a mushy layer.
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Superheating is a process in which a liquid is heated to a temperature above its boiling point without it actually boiling. This phenomenon occurs under certain conditions where the liquid is kept in a stable state, often due to the absence of nucleation sites or impurities that would facilitate the formation of bubbles. In practical applications, superheating is commonly observed in heating water or other liquids in a microwave.