What is chemical equilibrium?
Most reactions run both ways at once. Forward molecules combine into products while product molecules split back, and when the two rates exactly cancel, the concentrations freeze at their equilibrium values — not because anything stopped, but because every change is being undone as fast as it happens. That frozen ratio is captured in one number, the equilibrium constant K: products raised to their coefficients, divided by reactants raised to theirs. For the Haber reaction N2 + 3H2 = 2NH3, Kc = [NH3]2/([N2][H2]3). K is a tug-of-war score: much larger than 1 and products win, much smaller and reactants do.
The one rule to remember: only temperature changes K. Concentration, pressure and catalysts can shift where the equilibrium sits — they never move the constant itself.
Kc, Kp, and the bridge between them
When everything is dissolved, K uses concentrations (Kc); for gases, partial pressures work too (Kp). They are the same constant in different clothes, related by Kp = Kc(RT)Δn, where Δn counts gas molecules on the product side minus the reactant side. For the Haber reaction Δn = 2−4 = −2, so Kp is far smaller than Kc: with Kc = 4.0 at 500 K the sample computes Kp = 2.4×10−3. Thermodynamics closes the loop from the other side: ΔG° = −RT ln K, so ΔG° = −10 kJ/mol converts to K ≈ 57 (products favored) in the Gibbs panel — a small free-energy advantage is already a big ratio.
ICE tables: bookkeeping the approach
To predict where a reaction lands, lay out three rows: Initial concentrations, Change as ±coefficients·x, Equilibrium as their sum. Then solve for x. The ICE sample starts N2 + 3H2 at 1.0 and 3.0 M with no ammonia and Kc = 4.0; the solver finds x ≈ 0.65, so at equilibrium roughly 0.35 M N2, 1.06 M H2 and 1.29 M NH3 coexist — and plugging those back reproduces Kc = 4.0, which is exactly the verification the panel prints. Textbooks allow the “x is small” shortcut when less than 5% reacts; this page instead bisects the full equation, so no such assumption is needed.
Q vs K: which way from here?
The reaction quotient Q is K's formula applied to whatever concentrations exist right now, before equilibrium. Comparing them gives the direction: Q < K, forward; Q > K, reverse; Q = K, already there. The Q sample is a dramatic case — with Kc = 4.0 but Q = 4×104, the mixture is drowning in products and the reaction runs backward hard. Q is how a chemist predicts the direction of a proposed mixture before mixing it, and how the body decides, for instance, whether hemoglobin will pick up or release oxygen.
Le Chatelier: the stress rule
Disturb an equilibrium and it shifts to relieve the disturbance. Add reactant → it eats some; remove product → more forms; compress a gas mixture → it shifts toward the side with fewer gas molecules; heat it → it shifts in the endothermic direction. Note the fine print hiding in the key box above: temperature is the only stress that changes K itself (it re-rates the whole tug-of-war), while pressure and concentration only re-position it. A catalyst changes neither — it speeds forward and reverse equally, so equilibrium arrives faster but at exactly the same place. The Haber process is the industrial showcase: 4 gas molecules become 2, so high pressure pushes toward NH3, while temperature trades yield against rate.
Common misconceptions
- Equilibrium means the reaction stopped. Both directions keep running at full speed; only the net change is zero.
- Adding reactant raises K. It shifts the position forward and Q temporarily, then the system returns to the same K. Only temperature moves K.
- A catalyst increases the yield. It changes the time to reach equilibrium, never its location. Industrial yields move via temperature, pressure and continuous product removal.
- At equilibrium, reactant and product concentrations are equal. Equal concentrations mean K = 1, nothing more. K = 4.0 in the sample: products already four-times favored, yet plainly not equal.
Related tools: Acid-Base & pH (Ka and Kb are equilibrium constants), Thermochemistry (where ΔG and K come from), and Kinetics (the rates that set how fast equilibrium is reached).
