![]() ![]() The reason that the freezing of water does not violate the second law is because even though the system (ice) becomes more ordered and has lower entropy, the energy that is released to the surroundings makes those molecules move faster, which leads to an increase in the entropy of the surroundings. Now we can see the solution to our thermodynamic problem. That is, the freezing of water is an exothermic process. Because energy is conserved, this energy must be released to the surroundings as thermal (kinetic) energy. As we have seen previously, the formation of stabilizing interactions lowers the potential energy of the system. Consider when water freezes, the water molecules form stable interactions (hydrogen bonds). However for the universe as a whole (or more easily defined, the system and its surroundings) total entropy must increase. How can we resolve this seeming paradox? The answer lies in the fact that for any system the entropy may indeed decrease - water freezing is an example of this phenomenon. ![]() The problem with this is that we are all well aware of changes where the entropy apparently decreases. The Second Law of Thermodynamics tells us that for every change that occurs, the entropy of the universe must increase. How can it be that a change in which the entropy of the system decreases (for example freezing of ice) can occur? Are we forced to conclude that things we know to happen are impossible according to the Second Law of Thermodynamics? ![]() We can calculate how entropies change for materials as they go from gas to liquid to solid, and as we have predicted they decrease. How do we think about entropy in these systems? Doesn’t a substance become more ordered as we move it from gas to liquid to solid? Clearly the entropy of a solid is lower than that of a liquid (and the entropy of a liquid is lower than that of a gas). Now let us return to the situation with solids, liquids, and gases. ![]()
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