0
$\begingroup$

In a phase change the temperature remains constant but I can't understand why. The answers I have read state that during phase changes energy is only used to break bonds between molecules rather than increasing kinetic energy. However why is this the case? Why couldn't the supplied energy be used solely to increase the kinetic energy of particles?

$\endgroup$
3
  • $\begingroup$ If you need to overcome potential energy barrier to reach a different state, it is obvious that all energy cannot be used to increase kinetic energy when that state is reached. Like at ionization of atoms or at escaping gravity of a star or a planet. $\endgroup$
    – Poutnik
    Nov 20, 2022 at 6:31
  • $\begingroup$ "Temperature" is essentially movement of molecules. So you are on the right track thinking kinetic energy. Phase change is an energy sink moving from solid, to liquid to gas, so more energy input is required. In other words, temperature cannot rise because the molecules cannot move faster (increase KE) until phase change is completed. $\endgroup$ Nov 21, 2022 at 7:48
  • $\begingroup$ There are other phase changes like iron going from ferrite ( body centered cubic) to austenite ( face centered cubic). Thermocouples in a commercial heat-treat furnace would show a temperature drop while heating a load of steel as the phase change absorbed heat. Impressive to witness. $\endgroup$ Nov 21, 2022 at 21:03

2 Answers 2

3
$\begingroup$

It's easier to see in a specific example, e.g., boiling water. At 100°C and 1 atmosphere pressure, as more heat energy is added, the pressure of the water at the spot where heat is applied exceeds the pressure pushing on it, and bubbles form and rapidly expand, pulling away that added energy. The water vapor above the boiling water can be hotter than 100°C, but the liquid cannot... except for superheating.

Your presumption is not wrong: fairly clean water, without air bubbles, can be superheated well above 100°C, if left undisturbed -- a metastable state. The slightest disturbance, such as vibration, can cause the water to suddenly boil violently, "exploding" from the container. The remaining water is now below 100°C.

A similar phenomena is supercooling of water, in which heat energy has been transferred out of the system, yet the water does not freeze until it's disturbed (or is sufficiently cold, about −48 °C, below the freezing point of mercury). If you want to repeat the supercooling demonstration (much safer than superheating!), use a clean, smooth, plastic container and water that is clean, and has air bubbles removed by first boiling and allowing it to cool in a container with a lid (to keep out dust) before pouring it into the bottle.

However, those answers were also correct... once equilibrium has been reached. Eventually, superheated water will flash to steam. Supercooled water will eventually freeze. And even glass can crystallize. Don't let this dichotomy faze you, though.

$\endgroup$
3
$\begingroup$

This is one of the over simplifications that no one thinks about after being told. For a one component system the phase rule, F = C-P+2, states for two phases in equilibrium there is one degree of freedom. When evaporating a liquid, maintaining a constant P[vapor pressure] fixes the temperature of evaporation. A variation in P causes a variation in Temperature. The effects in melting are the same. To maintain a constant steady state Pressure an external energy source is required. If the energy source is the phase itself [or the heat source is insufficient] the temperature will change until equilibrium is reached; then the constant pressure means an invariant temperature.

In normal boiling, exposed to the atmosphere when the vapor pressure reaches atmospheric pressure P cannot increase since increasing evaporation simply pushes back the atmosphere. The Vapor pressure becomes constant and the temperature becomes constant. This is assuming sufficient energy input and not excessive heating and reasonable mixing and good exposure to the atmosphere.

$\endgroup$
3
  • 1
    $\begingroup$ So you can cause a phase change through compression or expansion, during which the system will stay on the coexistence line. But what do you mean by this: "then the constant pressure means an invariant temperature" ? $\endgroup$
    – Buck Thorn
    Nov 21, 2022 at 14:06
  • 1
    $\begingroup$ Just that! when the vapor pressure equals the constant atmospheric pressure the last freedom is removed and the temperature is constant. $\endgroup$
    – jimchmst
    Nov 21, 2022 at 20:23
  • 1
    $\begingroup$ If the vapor pressure is changed adiabatically the vapor will condense or liquid evaporate absorbing or releasing the heat of vaporization and the equilibrium point will change temperature. If the heat is supplied or removed externally the amount of the phases will change and P will remain constant. $\endgroup$
    – jimchmst
    Nov 21, 2022 at 20:27

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service and acknowledge you have read our privacy policy.

Not the answer you're looking for? Browse other questions tagged or ask your own question.