Calculate K_eq for the acid–base reaction shown. Which side is favored and why?

Acid-base equilibrium refers to the balance between the reactants and products in an acid-base reaction. It is characterized by the transfer of protons (H+) from the acid to the base. The equilibrium constant, K_eq, quantifies this balance, indicating the extent to which reactants are converted to products. A larger K_eq value suggests that the products are favored, while a smaller value indicates that the reactants are favored.

The strength of acids and bases is determined by their ability to donate or accept protons. Strong acids completely dissociate in solution, while weak acids only partially dissociate. Similarly, strong bases readily accept protons, whereas weak bases do so less effectively. Understanding the relative strengths of the acids and bases involved in the reaction is crucial for predicting the direction of the equilibrium and determining which side is favored.

Le Chatelier's Principle states that if a system at equilibrium is disturbed, the system will adjust to counteract the disturbance and restore a new equilibrium. In the context of acid-base reactions, if the concentration of either reactants or products changes, the equilibrium will shift towards the side that reduces the effect of that change. This principle helps in predicting how changes in concentration, pressure, or temperature will affect the position of equilibrium in acid-base reactions.