Question

During a phase change from solid to liquid, why does the internal energy of the system increase while the temperature remains constant according to the First Law of Thermodynamics?

Accepted Answer

According to the First Law of Thermodynamics, the change in internal energy of a system is equal to the heat added to it minus the work done by the system. During a phase change, such as melting, the substance is in a state of thermal equilibrium, where the incoming thermal energy is used exclusively to break the intermolecular forces—the attractive bonds holding particles in a rigid, fixed lattice structure—rather than increasing the kinetic energy, which is the energy of motion that determines temperature. Because temperature is a direct measure of the average kinetic energy of the particles, the temperature remains constant as long as the substance is transitioning between phases. Think of this like using a credit card to pay off an old debt; even though you are spending money, your &#x27;balance&#x27; (the temperature) stays the same because the energy is being consumed to settle the existing bond &#x27;debt&#x27; instead of being used to grow your savings. Consequently, the internal energy increases because the system has absorbed more heat, which is now stored as potential energy within the newly loosened arrangement of the particles.

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During a phase change from solid to liquid, why does the internal energy of the system increase while the temperature remains constant according to the First Law of Thermodynamics?