Left: molecular view (red=A, blue=B, purple=C). Right: concentration bars converging to equilibrium. Q/K meter shows reaction direction.
Chemical equilibrium occurs in a reversible reaction when the forward and reverse rates become equal, producing a stable mixture of reactants and products. The equilibrium constant K_eq — the ratio of product concentrations to reactant concentrations, each raised to their stoichiometric power — determines how far the reaction proceeds. Le Chatelier's principle states that if a system at equilibrium is disturbed, it will shift in the direction that opposes the disturbance and re-establishes equilibrium. This principle governs industrial processes such as the Haber-Bosch synthesis of ammonia (N₂ + 3H₂ ⇌ 2NH₃) and the production of sulfuric acid via the Contact process.
This simulator models the A + B ⇌ C reaction, letting you set initial concentrations of A and B, the equilibrium constant K_eq, and temperature. Buttons apply Le Chatelier perturbations — adding a reactant, removing a product, heating, or cooling — and the concentration-vs-time graph shows the system relaxing to its new equilibrium. The Q/K meter indicates whether the reaction is currently proceeding forwards (Q < K) or in reverse (Q > K).
What is the equilibrium constant K_eq?
K_eq is a dimensionless ratio defined at equilibrium as the product concentrations divided by the reactant concentrations, each raised to their stoichiometric coefficients. For A + B ⇌ C, K_eq = [C] / ([A][B]). A large K_eq (≫ 1) means the reaction strongly favours products; a small K_eq (≪ 1) means reactants predominate. K_eq depends only on temperature, not on initial concentrations or pressure.
What is the reaction quotient Q and why does it matter?
The reaction quotient Q has the same expression as K_eq but uses current (non-equilibrium) concentrations instead of equilibrium ones. Comparing Q to K tells you which way the reaction must shift: if Q < K, the forward reaction proceeds to make more product; if Q > K, the reverse reaction dominates; if Q = K, the system is at equilibrium. The Q/K meter in this simulator makes this comparison visual.
What does Le Chatelier's principle predict when you add more reactant A?
Adding reactant A increases [A], raising the denominator of Q and making Q < K. The system responds by shifting the equilibrium to the right — consuming some of the added A and B to produce more C — until Q returns to K. The concentrations of all three species change, but K_eq itself remains unchanged because temperature has not changed.
For an exothermic reaction (negative ΔH), heating shifts the equilibrium towards reactants and decreases K_eq, because the system responds by absorbing the added heat through the endothermic reverse reaction. For an endothermic reaction (positive ΔH), heating increases K_eq and shifts equilibrium towards products. The quantitative relationship is given by the van 't Hoff equation: d(ln K)/dT = ΔH/(RT²). In the Haber process, a compromise temperature (~450°C) is chosen to balance a high enough rate with a still-reasonable K_eq.
Removing C lowers [C], so Q = [C]/([A][B]) drops below K. The system re-establishes equilibrium by producing more C — shifting the reaction right. This is exploited industrially: in the Haber process, ammonia is continuously removed (liquefied) from the reactor to drive the reaction forwards despite the moderately low K_eq at operating temperatures.
Increasing pressure shifts equilibrium towards the side with fewer moles of gas. For N₂ + 3H₂ ⇌ 2NH₃, the left side has 4 moles of gas and the right has 2, so higher pressure favours ammonia production. For reactions with equal moles of gas on both sides (like H₂ + I₂ ⇌ 2HI), pressure has no effect on equilibrium position. Pressure does not affect K_eq for ideal gases, but it does change concentrations and thus Q.
The ICE table (Initial, Change, Equilibrium) is a systematic way to find equilibrium concentrations. You list the initial concentrations, express the change in terms of a variable x (the extent of reaction), write the equilibrium expressions as (initial ± x), and substitute into the K_eq equation to solve for x. For the A + B ⇌ C model in this simulator, the quadratic formula solves the resulting equation analytically.
No. A catalyst speeds up both the forward and reverse reactions equally, so equilibrium is reached faster but the final equilibrium concentrations and K_eq are unchanged. In the Haber process, iron catalyst is used to achieve equilibrium at practical timescales (milliseconds rather than years) at 450°C — without the catalyst, the reaction is effectively frozen at that temperature.
A small K_eq means the equilibrium lies far to the left (reactant side), so the conversion percentage — the fraction of limiting reactant converted to product — is low. For example, with K_eq = 0.01 and equal initial concentrations, the equilibrium conversion may be only a few percent. Industrially, low-conversion reactions often employ recycling: unreacted feedstock is separated and returned to the reactor.
Enzyme-catalysed reactions in biology obey the same equilibrium laws. The Michaelis-Menten model of enzyme kinetics is built on a rapid equilibrium between enzyme (E) and substrate (S) forming a complex (ES), before the product (P) is released. Le Chatelier-like effects appear when metabolic products are continuously consumed by downstream reactions, keeping Q ≪ K and driving biosynthesis forward efficiently.