Which statement correctly describes the relationship between Q (reaction quotient) and K (equilibrium constant) in predicting the direction of a reaction?

Which statement correctly describes the relationship between Q (reaction quotient) and K (equilibrium constant) in predicting the direction of a reaction?

• A. If Q < K, the system shifts to the left.
• B. If Q > K, the system shifts to the right.
• C. If Q > K, the system shifts to the left.
• D. If Q = K, the system is not at equilibrium.

Answer: (C)

The key idea is comparing the current state, Q, to the equilibrium constant, K. If Q is greater than K, there are too many products compared with reactants, so the system shifts toward the reactants to restore balance. That means a shift to the left. Conversely, if Q is less than K, the system shifts toward more products (to the right). When Q equals K, the system is at equilibrium with no net change. Since the statement describes Q > K leading to a leftward shift, it correctly captures how the reaction responds to excess products. Remember, K is fixed at a given temperature, while Q can change with concentrations or pressures.

The key idea is comparing the current state, Q, to the equilibrium constant, K. If Q is greater than K, there are too many products compared with reactants, so the system shifts toward the reactants to restore balance. That means a shift to the left. Conversely, if Q is less than K, the system shifts toward more products (to the right). When Q equals K, the system is at equilibrium with no net change. Since the statement describes Q > K leading to a leftward shift, it correctly captures how the reaction responds to excess products. Remember, K is fixed at a given temperature, while Q can change with concentrations or pressures.