The best example of decreasing entropy is water freezing into ice. When liquid water molecules, which move freely and occupy many possible arrangements, lock into the rigid, repeating lattice of an ice crystal, the system becomes dramatically more ordered. Its entropy drops. Crystallization is the textbook answer to this question because the change is vivid, measurable, and easy to visualize: a disordered liquid transforms into a highly structured solid.
But that answer only makes sense once you understand what “decreasing entropy” really means and why it doesn’t violate the laws of physics. Several natural processes decrease entropy locally, and comparing them helps you see why crystallization stands out as the clearest example.
What Entropy Actually Measures
Entropy is a measure of how many different microscopic arrangements a system’s particles can take while still looking the same from the outside. A gas has high entropy because its molecules can be almost anywhere, moving in any direction. A crystal has low entropy because each molecule sits in a fixed position within a repeating pattern. When a system shifts from more possible arrangements to fewer, its entropy decreases.
The second law of thermodynamics says the total entropy of the universe always increases. Entropy is generated everywhere and always, at every scale, and it cannot be destroyed by any means. But entropy can decrease locally. The key distinction is that a local decrease in entropy requires energy transfer to the surroundings, which generates even more entropy elsewhere. The freezer that turns your water into ice cubes pumps heat into your kitchen, increasing the entropy of the room more than the ice’s entropy went down.
Why Crystallization Is the Strongest Example
When a liquid crystallizes, molecules go from a relatively disordered state to a highly ordered lattice. The entropy change is negative and substantial. For protein crystals (a well-studied case), the non-electrostatic entropy of crystallization has been measured at roughly negative 96 joules per kelvin per mole. That’s a large, clean drop in disorder for the material itself.
Crystallization stands out as the best example for several reasons. The before-and-after contrast is stark: liquid molecules have many accessible positions and orientations, while crystal molecules have essentially one. The process is common and observable in everyday life, from ice forming on a pond to salt crystals appearing as water evaporates. And the entropy decrease happens in a single, well-defined step rather than through a long chain of complex reactions.
The trade-off is straightforward, too. When water freezes, it releases heat (the “heat of fusion”) into the surroundings. That released thermal energy increases the entropy of the environment by more than the crystal’s entropy decreased, keeping the second law intact.
Other Strong Examples of Decreasing Entropy
Crystallization isn’t the only process that reduces local entropy. Several others are worth understanding, especially because they appear on exams and in real-world discussions.
Protein Folding
An unfolded protein is a long, floppy chain of amino acids that can adopt an enormous number of shapes. When it folds into its functional three-dimensional structure, the chain collapses into one specific arrangement, and the protein’s own entropy plummets. The process works through a series of trade-offs: water-repelling side chains cluster together in the protein’s interior, which frees water molecules that had been forced into ordered shells around those side chains. The freed water molecules gain entropy, compensating for the entropy the protein lost. Internal hydrogen bonds then lock the structure in place. This is why some proteins never fold at all. When the energy trade-off doesn’t work out, the chain stays disordered.
Living Cells
A living cell is one of the most impressively low-entropy structures in nature. It keeps thousands of different molecular species separated into specific compartments: the nucleus, the mitochondria, the endoplasmic reticulum, and others. Maintaining this organization requires a constant supply of energy. Cells import chemical energy through metabolism and export heat and waste products, dumping entropy into the environment. Without that continuous energy flow, the cell’s internal order would collapse. This is, in a real sense, what death is: the moment a system stops exporting entropy and disorder takes over.
Earth’s Energy Balance
Earth itself is a large-scale entropy-decreasing system. The planet absorbs high-energy, low-entropy sunlight (photons packed with energy at roughly 5,500°C) and reradiates the same total energy back into space as lower-energy infrared radiation at much cooler temperatures. Because the outgoing radiation is cooler, it carries far more entropy than the incoming sunlight did. That difference, the gap between low-entropy input and high-entropy output, is what powers every weather system, ocean current, and living organism on the planet.
How to Identify Decreasing Entropy
If you’re answering a multiple-choice question, look for the option where matter goes from a more disordered state to a more ordered one. The classic signals are:
- Gas turning into a liquid or solid. Fewer accessible arrangements means lower entropy.
- Liquid turning into a solid. Crystallization is the most direct example.
- Dissolved particles forming a precipitate. Ions scattered through a solution lock into a crystal lattice.
- Fewer gas molecules after a reaction. If a chemical reaction produces fewer moles of gas than it started with, the entropy of the system typically decreases.
Processes that go the other direction, like ice melting, a solid dissolving, or a gas expanding, all increase entropy. The more freedom particles gain, the higher the entropy.
Why It Doesn’t Break the Second Law
The most common confusion around decreasing entropy is the belief that it shouldn’t be possible. The second law says entropy always increases, so how can anything become more ordered? The answer is that the second law applies to isolated systems, meaning systems that exchange neither energy nor matter with their surroundings. Almost nothing in daily life is truly isolated. Your freezer, a growing crystal, a living cell: all of these exchange energy with their environment.
When you account for both the system and its surroundings, the total entropy always goes up. The local decrease inside the system is always paid for by a larger increase somewhere else. Entropy can decrease locally, but it cannot be destroyed anywhere. Without this energy exchange with the environment, there would be no formation of cyclones, crystals, or life.
So when a question asks for “the best example of decreasing entropy,” crystallization (and specifically water freezing into ice) is the cleanest answer. It’s a single-step physical change with a large, measurable drop in disorder, a clear mechanism for how the entropy is exported, and an intuitive before-and-after comparison that makes the concept concrete.