Why a Phase Transition Temperature Remains Constant

A common misconception about phase transitions is that latent heat is used exclusively to break intermolecular bonds, leaving the kinetic energy and therefore the temperature unchanged. In reality, the release of molecules from fixed positions involves changes in kinetic energy. So why does the temperature remain constant? This paradox arises from a more fundamental misunderstanding of temperature itself. Contrary to conventional textbook definitions, our study suggests that temperature is not a direct measure of a system’s kinetic energy, but is instead more fundamentally linked to its radiative energy. Since radiative energy is emitted by potential energy components, such as intermolecular bonds, temperature is more closely tied to potential energy than previously recognized. This perspective is supported by infrared thermometers, which measure temperature based on radiated energy, not molecular motion. Further evidence comes from discrepancies in the specific heat of ideal gases: predictions based on the kinetic theory of gases consistently fall short of observed values, with the deficit corresponding to the number and type of molecular bonds. This implies that the unaccounted specific heat arises from increases in potential energy rather than kinetic energy. Accordingly, the constancy of temperature during a phase transition is not due to fixed kinetic energy, but to an upper limit on the system’s potential energy, determined by the strength of intermolecular bonds. These bonds, like springs, store increasing amounts of energy as they are stretched with rising temperatures. A phase transition occurs when the bonds reach their maximum extension and begin to rupture, driven by transimpacts—impacts caused by atomic electron transitions. The energy threshold of these transimpacts defines the maximum potential energy the system can sustain, which in turn determines its radiative output and thus fixes the temperature during the phase change.