Final answer:
The temperature remains constant during a phase change because energy is used to overcome intermolecular forces rather than to increase temperature. This is observed during melting and boiling of substances, such as water, where the temperature stays at the melting or boiling point respectively until the phase change is complete.
Step-by-step explanation:
The temperature does not increase during a phase change, even though energy is being absorbed by the system. This is because the absorbed energy is used to break the attractive forces between molecules rather than increase the kinetic energy, which determines the system's temperature. During a phase change such as melting or boiling, the process is isothermal, meaning the temperature remains constant while the phase change occurs. For instance, when ice melts, the temperature stays at 0°C until all the ice has turned to water. The same logic applies when water boils; it remains at 100°C as long as there's liquid water present, regardless of how much energy is added.
If heat is added at a constant rate, as shown in heat vs. temperature graphs of a substance undergoing a phase change, the horizontal parts of the graph where temperature does not change corresponds to the phase changes. The length of these horizontal segments is directly proportional to the amount of heat energy known as the latent heat, absorbed or released during the phase transition. The enthalpy of vaporization is typically larger than the enthalpy of fusion, which is why the horizontal segment during boiling is longer than that during melting.