7.3 Reaction Quotient and Equilibrium Constant

The Reaction Quotient, Q

The reaction quotient (Q) has the same mathematical form as the equilibrium constant, using the same expression of product concentrations over reactant concentrations (each raised to their stoichiometric coefficients). Unlike K, Q can be calculated at any point during a reaction, not just at equilibrium. Comparing Q to K tells you which direction a reaction must shift to reach equilibrium.

Comparing Q and K

If Q < K, there are too few products relative to reactants, so the reaction shifts forward (toward products). If Q > K, there are too many products, so it shifts reverse (toward reactants). If Q = K, the system is already at equilibrium and there's no net shift. This comparison predicts reaction direction without needing to solve for concentrations.

Equilibrium Constant Expression

For a reaction aA + bB ⇌ cC + dD, K = [C]^c[D]^d / [A]^a[B]^b. Pure solids and liquids are omitted from the expression since their activities are defined as 1. K is temperature-dependent but independent of initial concentrations or pressure changes — only temperature changes its value.

Key terms

Reaction Quotient (Q)

An expression with the same form as K, calculated at any point in a reaction to predict the direction of shift needed to reach equilibrium.

Q < K

Reaction shifts forward (toward products) because not enough product is present yet.

Q > K

Reaction shifts in reverse (toward reactants) because too much product is present.

Q = K

System is at equilibrium; no net shift occurs in either direction.

What changes the value of K?

Only a change in temperature changes K; concentration, pressure, and catalysts do not.

Try an AP-style question

What is K for N₂(g) + 3H₂(g) ⇌ 2NH₃(g)?

1. [NH₃] / ([N₂][H₂])
2. [N₂][H₂]³ / [NH₃]²
3. [NH₃]² / ([N₂][H₂]³)
4. 2[NH₃] / ([N₂] · 3[H₂])