Reaction Quotient and Equilibrium Constant

In a reaction at equilibrium, individual concentrations can look stuck but aren't fixed in any deep sense — add more of one species and the others shift. What stays constant is a specific ratio: products raised to their coefficients, divided by reactants raised to theirs. That ratio is K, the equilibrium constant, and it's fixed at a given temperature because it reflects the balance between forward and reverse reaction rates, not the amount of any one substance.

To build the expression, start from the balanced equation. For a A + b B <→ c C + d D, K = [C]^c[D]^d / ([A]^a[B]^b). For N₂(g) + 3 H₂(g) <→ 2 NH₃(g), that gives K = [NH₃]² / ([N₂][H₂]³). If a solid or pure liquid shows up, like CaCO₃(s) <→ CaO(s) + CO₂(g), leave it out: K = [CO₂]. A pure solid or liquid has a fixed density and molar volume, so its 'concentration' never changes — that constant gets folded into K itself rather than tracked as a variable.

Q, the reaction quotient, uses the exact same expression but can be calculated at any point in a reaction, not just at rest. Comparing Q to K tells you which way the reaction will shift: if Q < K there aren't enough products yet, so the reaction runs forward; if Q > K it runs in reverse; if Q = K the system is already at equilibrium. This comparison is the bridge into using ICE tables and predicting shifts from Le Chatelier's principle.