When energy is transferred to a substance, typically as heat, we often expect a rise in temperature. However, the amount of energy added is not always directly proportional to a change in the reading on a thermometer. Adding heat can result in two distinct outcomes: the substance gets hotter, or its physical state changes, such as melting or boiling. Understanding why energy transfer sometimes causes a temperature increase and sometimes causes a change of phase requires looking closely at what the energy is doing to the molecules.
When Energy Translates to Temperature Increase
When energy is introduced and the temperature rises, that energy is called sensible heat because the change can be “sensed” with a thermometer. This added energy increases the motion of the substance’s molecules and atoms. In a solid, particles vibrate more vigorously; in a liquid or gas, they move around and collide more quickly.
This increased speed and vibration is directly related to the substance’s average kinetic energy, which is the physical measurement we interpret as temperature. The relationship between energy input and temperature change is quantified by the material’s specific heat capacity. Specific heat is a unique property representing the amount of energy required to raise the temperature of a specific mass of that material by one degree. For instance, water has a high specific heat, meaning it takes a large amount of energy to raise its temperature compared to many other materials.
The Energy Required to Change State
Sometimes, a substance absorbs significant amounts of energy without any corresponding temperature increase, a phenomenon known as latent heat. This “hidden” energy is completely occupied with changing the substance’s physical state. During a phase transition, such as melting or boiling, the temperature remains constant.
The absorbed energy does not increase the average kinetic energy of the molecules. Instead, the energy is used to overcome and break the attractive forces, called intermolecular bonds, that hold the molecules together. For melting, this absorbed energy (latent heat of fusion) weakens the rigid structure of the solid, allowing molecules to move freely as a liquid. For boiling, the energy (latent heat of vaporization) separates the molecules into a gaseous state. Because the energy is consumed by changing the internal molecular arrangement and increasing the potential energy, the measured temperature does not change until the transition is complete.
How Energy Prioritizes Molecular Action
The reason an energy transfer does not always result in a phase change is that the energy prioritizes increasing the molecular motion, and thus the temperature, until a specific transition point is reached. When a substance is heated, the initial energy input raises the kinetic energy of its molecules, leading to a steady temperature increase. This continues until the substance reaches its melting or boiling point, a temperature unique to that material.
At this precise temperature, the energy input is diverted from raising the temperature to performing the work of a phase change. For example, if you heat ice below \(0^\circ\text{C}\), the temperature rises until it hits \(0^\circ\text{C}\). Once at \(0^\circ\text{C}\), any additional energy is entirely spent on breaking the bonds between the water molecules to convert the solid ice into liquid water, keeping the temperature fixed at \(0^\circ\text{C}\) until all the ice is gone.
This period of constant temperature is represented as a plateau on a heating curve, indicating that the potential energy of the molecules is increasing while their average kinetic energy is not. Once the entire mass has transitioned to the new state, the added energy reverts to increasing the molecular kinetic energy. The temperature begins to rise again until the next phase transition point is reached.