How is entropy change related to reaction reversibility?

Entropy change is related to reaction reversibility as reversible reactions tend to have an entropy change close to zero.

In more detail, entropy is a measure of the disorder or randomness of a system. In a chemical reaction, the entropy change (∆S) is the difference in entropy between the products and the reactants. If the entropy of the products is greater than the entropy of the reactants, the entropy change is positive, indicating an increase in disorder. Conversely, if the entropy of the products is less than the entropy of the reactants, the entropy change is negative, indicating a decrease in disorder.

Reversible reactions are those that can proceed in both the forward and reverse directions. In a reversible reaction at equilibrium, the rates of the forward and reverse reactions are equal, and the concentrations of the reactants and products remain constant. This means that the system is in a state of maximum disorder, or maximum entropy.

The second law of thermodynamics states that the total entropy of an isolated system can never decrease over time, and is constant if and only if all processes are reversible. Therefore, for a reaction to be reversible, the entropy change must be close to zero. This is because a large positive entropy change would mean that the reaction proceeds almost entirely in the forward direction, while a large negative entropy change would mean that the reaction proceeds almost entirely in the reverse direction. In both cases, the reaction would not be reversible.

In practice, however, no reaction is perfectly reversible because some energy is always lost to the surroundings as heat, causing an increase in the entropy of the surroundings. This is why the entropy change for a reversible reaction is close to zero, but not exactly zero. The concept of entropy change and its relationship to reaction reversibility is a fundamental aspect of thermodynamics, and is crucial for understanding many chemical reactions and processes.

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