How do endothermic and exothermic reactions differ in terms of equilibrium shifts?

Endothermic and exothermic reactions differ in how temperature changes shift the equilibrium position.

In an exothermic reaction, heat is released, making it a product of the reaction. According to Le Chatelier's principle, if you increase the temperature (add more product), the equilibrium will shift to the left, favouring the reactants. Conversely, if you decrease the temperature (remove product), the equilibrium will shift to the right, favouring the products. This is because the system is trying to counteract the change and re-establish equilibrium.

On the other hand, in an endothermic reaction, heat is absorbed, making it a reactant. If you increase the temperature (add more reactant), the equilibrium will shift to the right, favouring the products. If you decrease the temperature (remove reactant), the equilibrium will shift to the left, favouring the reactants. Again, this is the system's way of counteracting the change to re-establish equilibrium.

It's important to remember that these shifts in equilibrium do not change the equilibrium constant, K, at a given temperature. The value of K only changes if the temperature changes. This is because K is dependent on the standard Gibbs free energy change, ΔG°, which is temperature dependent.

In summary, the direction of an equilibrium shift due to a temperature change depends on whether the reaction is endothermic or exothermic. For exothermic reactions, increasing temperature shifts the equilibrium to the left, while decreasing temperature shifts it to the right. For endothermic reactions, increasing temperature shifts the equilibrium to the right, while decreasing temperature shifts it to the left. This is all in accordance with Le Chatelier's principle, which states that a system at equilibrium will respond to a change in conditions in a way that counteracts the change.