Gibbs free energy determines whether a reaction is spontaneous from "summary" of Chemistry by Russell Kuhtz
The Gibbs free energy is a crucial concept in determining whether a reaction is spontaneous or not. When we talk about spontaneity in chemistry, we are referring to whether a reaction will occur without any external influence. In other words, will the reaction happen on its own without needing any extra push?To understand this concept, we need to look at the equation for Gibbs free energy, which is ΔG = ΔH - TΔS. In this equation, ΔG represents the change in Gibbs free energy, ΔH is the change in enthalpy, T is the temperature in Kelvin, and ΔS is the change in entropy. If the change in Gibbs free energy (ΔG) for a reaction is negative, it means that the reaction is spontaneous. This is because a negative ΔG indicates that the products of the reaction are in a lower energy state than the reactants. In simpler terms, the reaction will naturally move towards the products because they are more stable. Conversely, if the change in Gibbs free energy (ΔG) for a reaction is positive, it means that the reaction is non-spontaneous. In this case, the products are in a higher energy state than the reactants, so the reaction will not proceed on its own. It would require an input of energy to overcome this energy barrier and make the reaction happen. The temperature also plays a crucial role in determining the spontaneity of a reaction. A reaction that is non-spontaneous at a low temperature may become spontaneous at a higher temperature if the change in entropy (ΔS) is positive. This is because the TΔS term in the Gibbs free energy equation can outweigh the positive ΔG, making the overall ΔG negative and the reaction spontaneous. Therefore, by calculating the change in Gibbs free energy for a reaction, we can predict whether the reaction will occur spontaneously or not. This is a fundamental concept in chemistry that helps us understand the behavior of chemical reactions and the factors that influence their spontaneity.
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