Three physical changes are endothermic: melting, vaporization, and sublimation. Each one moves matter from a more ordered state to a less ordered one (solid to liquid, liquid to gas, or solid straight to gas), and that transition requires an input of energy. The reverse processes, freezing, condensation, and deposition, release energy and are exothermic.
Why These Changes Absorb Energy
Molecules in a solid are held tightly together by attractive forces between them. To loosen those molecules into a liquid, or free them entirely into a gas, energy has to come in from the surroundings to pull the molecules apart. That absorbed energy doesn’t raise the temperature of the substance while the change is happening. Instead, it goes entirely toward breaking those intermolecular bonds. This is why a pot of water stays at 100 °C the entire time it boils, even though you’re continuously adding heat from the stove. The energy is being consumed by the phase change itself.
Melting (Fusion)
Melting is the transition from solid to liquid. Ice turning into water is the most familiar example. It takes about 6 kJ of energy to melt one mole (roughly 18 grams) of ice at 0 °C. That may sound modest, but it’s why a glass of ice water stays cold for so long: the ice absorbs a significant amount of heat from the surrounding liquid before it fully melts. Other everyday examples include butter softening on a warm counter and candle wax liquefying near a flame.
Vaporization (Boiling and Evaporation)
Vaporization is the transition from liquid to gas. It happens at the boiling point when you heat a liquid, but it also happens at lower temperatures through evaporation at the surface. Both are endothermic. Vaporization demands far more energy than melting because gas molecules must separate almost completely from one another. For water, the energy required to vaporize one mole is about 40.7 kJ, nearly seven times the energy needed to melt the same amount of ice. To put it in practical terms, turning 100 grams of boiling water entirely into steam takes about 226 kJ of energy.
This is the principle behind one of your body’s most important cooling systems. When you sweat, the liquid on your skin absorbs heat from your body as it evaporates. The highest-energy water molecules escape into the air, carrying that thermal energy with them and leaving your skin cooler. This evaporative cooling is the primary way humans shed excess heat during exercise or in hot environments, and it works precisely because vaporization is endothermic.
Sublimation
Sublimation is the jump from solid directly to gas, skipping the liquid phase entirely. The most common example is dry ice, which is solid carbon dioxide. At normal atmospheric pressure, it transitions straight from a solid at -78.5 °C into carbon dioxide gas, which is why it produces that dramatic fog without ever forming a puddle. Snow and ice also sublimate in cold, dry conditions. On a sunny winter day with low humidity, snow can slowly disappear into water vapor without melting first. Because sublimation essentially combines the energy costs of both melting and vaporization into a single step, it absorbs the most energy of the three endothermic phase changes.
Endothermic Dissolving
Dissolving a solid in a liquid is also a physical change, and some dissolving processes are endothermic. The clearest example is the instant cold pack you might find in a first-aid kit. These packs contain a salt (commonly ammonium nitrate) and a pouch of water. When you squeeze the pack and the salt mixes with the water, the dissolving process absorbs so much heat from the surroundings that the pack becomes ice cold within seconds. The temperature drop can be dramatic enough to freeze a thin layer of water on the outside of the container. Not every dissolving process is endothermic, though. Table salt dissolving in water, for instance, absorbs very little heat, and some substances release heat when they dissolve.
How to Tell a Process Is Endothermic
The simplest way to recognize an endothermic physical change in everyday life is to notice whether the surroundings get cooler. A cold pack chills your hand. A puddle evaporating on a hot sidewalk makes the concrete slightly cooler beneath it. Ice melting in a drink pulls warmth from the liquid around it. In each case, energy flows from the environment into the substance undergoing the change.
In chemistry, this is expressed as a positive enthalpy change. A positive value means the system absorbed energy. Melting, vaporization, and sublimation all carry positive enthalpy values, confirming they are endothermic. Their reverse processes carry negative values, confirming they release energy. This sign convention is a reliable shortcut: if the enthalpy change is positive, the physical change is endothermic.