What is the main difference between K and Q in thermodynamics?

The primary distinction between K (the equilibrium constant) and Q (the reaction quotient) lies in their definitions and conditions of application. K is specifically determined at equilibrium for a particular reaction at a given temperature, meaning it reflects the ratio of concentrations (or partial pressures) of products to reactants when the system has reached a state of balance where the rates of the forward and reverse reactions are equal.

On the other hand, Q can be calculated at any point during the reaction, regardless of whether the system is at equilibrium or not. This allows Q to provide insight into the current state of the reaction relative to equilibrium. By comparing Q to K, one can predict the direction in which the reaction will shift to reach equilibrium: if Q is less than K, the reaction will proceed forward to produce more products; if Q is greater than K, the reaction will shift backward to produce more reactants.

This foundational understanding of Q and K is crucial for predicting the behavior of chemical reactions under various conditions.