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roam
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To show that entropy is not the same thing as disorder (what people intuitively accept as disorder) my textbook gives an example of crystallization in a supersaturated solution. And it argues that since both temprature and disorder decrease the entropy must decrease also, but it does not. Hence giving a contradiction to the disorder interpretation:
Unlike what the book says, in reality the entropy must decrease because no such system is truly isolated. Right?
Also, doesn't the 2nd law say that entropy only tends to increase in an isolated system (it can decrease locally within an isolated system)? So, wouldn't stacking some coins would have sufficed as example? So why give this example in particular?
Consider an isolated supersaturated solution a liquid in which a solid has been dissolved to a concentration greater than it would be for equilibrium. Such a solution is unstable. A crystal suddenly and spontaneously forms in the solution. The entropy of the system cannot decrease. Yet the appearance of the crystal certainly would be regarded as an increase in order. But in this example the temperature of the system could decrease. How on Earth can we retain the disorder interpretation of entropy when the system has undergone a partial transition from liquid to solid and its temperature has also decreased?
Unlike what the book says, in reality the entropy must decrease because no such system is truly isolated. Right?
Also, doesn't the 2nd law say that entropy only tends to increase in an isolated system (it can decrease locally within an isolated system)? So, wouldn't stacking some coins would have sufficed as example? So why give this example in particular?