Reaction Quotient Q (aA + bB ⇌ cC)

Q is built from exactly the same expression as the equilibrium constant — products over reactants, each raised to its balanced coefficient — but you are allowed to evaluate it with any concentrations, not just equilibrium ones. That is the whole point. Q is a snapshot of where a mixture currently stands; K is where it is heading. Compare them and you have the direction of shift for free: Q < K means too much reactant, so the reaction runs forward; Q > K means too much product, so it reverses; Q = K means the mixture is already at equilibrium and nothing net happens.

Each concentration is divided by the standard state c° = 1 mol/L before it is raised to its power, which is what makes Q a pure number even when the powers do not balance. This page covers the very common aA + bB ⇌ cC shape; for a reaction with two products and matching one-to-one coefficients, use the Kc calculator instead. A species that does not appear on the reactant side can be switched off by entering a coefficient of 0 — pure solids and pure liquids are always left out, which is why the water in an aqueous equilibrium never shows up.

Henry Le Chatelier, whose 1884 principle is the qualitative version of this comparison, paid dearly for not having the quantitative one. He ran nitrogen and hydrogen together at high pressure looking for ammonia, an explosion in his laboratory nearly killed an assistant, and he abandoned the work — later calling it "the greatest blunder of my scientific career". The explosion was caused by air left in the apparatus, not by the equilibrium. Fritz Haber took it up two decades later. Worked case, on Le Chatelier's own reaction N₂ + 3H₂ ⇌ 2NH₃: with [N₂] = 0.50 M, [H₂] = 0.20 M and [NH₃] = 0.10 M, Q = (0.10)²/[(0.50)(0.20)³] = 0.010/0.0040 = 2.5. If K at that temperature is 0.060, then Q > K and the mixture will run backwards, decomposing ammonia — the difference between a plant that makes fertiliser and one that does not.