During phase change in matter, why doesn't the temperature change?

I was working on something in school and came across the question:

Why does the temperature not change much during a phase change?

I'm really not sure why this happens in matter and I couldn't find an answer in my school resources. Does anyone here know?

...

So, how could there be a change in heat during a state change without a change in temperature?

"During a change in state the heat energy is used to change the bonding between the molecules. In the case of melting, added energy is used to break the bonds between the molecules. In the case of freezing, energy is subtracted as the molecules bond to one another. These energy exchanges are not changes in kinetic energy. They are changes in bonding energy between the molecules.

"If heat is coming into a substance during a phase change, then this energy is used to break the bonds between the molecules of the substance. The example we will use here is ice melting into water. Immediately after the molecular bonds in the ice are broken the molecules are moving (vibrating) at the same average speed as before, so their average kinetic energy remains the same, and, thus, their Kelvin temperature remains the same."

• Couldn't we say that during the change the molecules absorb the heat as kinetic energy, move faster so they can overcome the attractions (potential energy)? But the attraction are still there. So after a little time the molecules would be in a greater distance (higher potential energy) with same kinetic (conservation of energy) so the temperature will remain the same. Oct 29 '20 at 11:49

For a first-order phase transition, you need to add the enthalpy of the phase transition. As an example, starting with ice below the melting point, you pump heat in, and raise the temperature. When you hit the melting temperature, the heat you put in goes towards the enthalpy of melting, and starts converting ice (sold) to water (liquid). Additional heat continues to melt more of the ice. Once all of the ice is converted (still at the melt temperature), than more heat starts increasing the temperature of the water.

A second-order phase transition does not have an enthalpy associated with it.

• Thanks so much for your answer. This seems to tentatively make sense on an initial read, but unfortunately I am not familiar with the concept of enthalpy and how it effects the phase transitions. Would you mind explaining this in layman's terms inline with your succinct explanation?
– GDP2
Sep 2 '14 at 19:14
• @Ylluminarious - The concepts of enthalpy (H) and entropy (S) are fundamental to thermodynamics. The Gibbs free energy $G = H - TS$. A first order phase transition has a discontinuity in $G$, which is the enthalpy of the phase transition. An introductory thermodynamics textbook should handle this pretty early on. Sep 2 '14 at 19:21
• @Ylluminarious perhaps a simple introductory video such as this could help clarify the concept of enthalpy. Sep 2 '14 at 19:49

It is because the heat energy is used to overcome the inter-molecular force of attraction but when this inter-molecular force is broken it changes it state and the temperature starts increasing.

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Mix water and ice in a container. Keep stirring the mix while the ice is melting. Then the following three statements are true:

1. The temperature of the water is greater than or equal to 0.

2. The temperature of the ice is less than or equal to 0.

3. The water and the ice have roughly the same temperature.

It follows that the temperature of the mix is roughly 0.

• Statements 1 and 2 are only true under certain conditions. See physics.stackexchange.com/q/60170/12613. Statement 3 may not be true. If I drop ice at -5C into an equal mass of water at +50C it will take an appreciable time before the temperatures of the water and ice approach a common temperature. Sep 3 '14 at 13:54