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Every phase change — melting, freezing, boiling, condensation — involves energy being absorbed or released. Melting and boiling are endothermic: they require energy input to overcome the attractions between particles. Freezing and condensation are exothermic: they release energy as particles form stronger attractions. During a phase change, the temperature of the substance stays constant even though energy is being added or removed — all the energy goes into changing the arrangement of particles rather than increasing their speed.
Heat ice in a beaker and record the temperature every minute. Graph the results (time vs. temperature) to create a heating curve. The flat sections during melting (0°C) and boiling (100°C) dramatically illustrate that temperature does not change during phase transitions, even though you are constantly adding heat.
You already know that substances can change between solid, liquid, and gas phases, and that reactions can be endothermic or exothermic. Now it is time to connect these two ideas: every phase change involves a transfer of energy, and understanding that energy transfer explains some surprising behaviors.
When you heat a block of ice, its temperature rises steadily — until it reaches 0°C. At that point, something unexpected happens: the temperature stops rising. You keep adding heat, but the thermometer stays at 0°C. The ice is melting — transforming from solid to liquid — and all the energy you are adding is going into breaking the attractions between water molecules in the ice structure. This energy is not making the particles move faster (which would increase temperature); it is pulling them apart from their rigid arrangement. Only after all the ice has melted does the temperature start climbing again. Melting is endothermic — it absorbs energy.
The same thing happens at 100°C during boiling. The temperature plateaus again as all the incoming energy goes into breaking the attractions between liquid water molecules, launching them into the gas phase. This is why a pot of boiling water stays at 100°C no matter how high you turn the flame — the extra energy just makes the water boil faster, not hotter. Boiling is endothermic too.
If you graphed temperature versus time as you heated ice all the way to steam, you would see a heating curve — a line that rises, flattens, rises, flattens, and rises again. The flat sections correspond to the two phase changes (melting and boiling), where temperature holds steady while energy transforms the substance's structure.
The reverse processes — freezing and condensation — are exothermic. When liquid water freezes, the molecules slow down and lock into the organized structure of ice, releasing energy into the surroundings. This might seem strange, but it explains real phenomena. Fruit growers sometimes spray water on crops during cold snaps — when the water freezes, it releases just enough heat to protect the fruit from frost damage. Condensation releases energy too, which is why humid air feels warmer than dry air at the same temperature, and why steam burns are particularly severe: the steam releases extra energy when it condenses on your skin.
The key principle is that energy and particle arrangement are connected. Adding energy breaks attractions between particles (solid → liquid → gas). Removing energy allows attractions to form (gas → liquid → solid). Temperature only changes when the energy is speeding up or slowing down particles — during a phase change, the energy is all going into restructuring the arrangement, and the temperature holds steady.