Equilibrium Constant (Keq) Calculator

The equilibrium constant (Keq) describes the ratio of product concentrations to reactant concentrations at equilibrium.

For the reaction: aA + bB ⇌ cC + dD

Keq = Kc = [C]^c × [D]^d / ([A]^a × [B]^b)

All concentrations are in mol/L, measured at equilibrium.

Rules:

• Pure solids and pure liquids are excluded from the expression (their “concentration” is constant)
• Keq is temperature-dependent but pressure/concentration independent (at equilibrium)
• Large Keq (» 1): equilibrium favors products
• Small Keq (« 1): equilibrium favors reactants
• Keq ≈ 1: significant amounts of both reactants and products present

Kp for gas-phase reactions:

Kp = Kc × (RT)^Δn

Where Δn = (moles of gaseous products) − (moles of gaseous reactants). R = 0.08206 L·atm/mol·K.

Relationship to ΔG°:

ΔG° = −RT ln Keq

Temperature effect on Keq: Using van’t Hoff equation:

d(ln Keq)/dT = ΔH° / RT²

For exothermic reactions (ΔH° < 0): Keq decreases with temperature. For endothermic reactions (ΔH° > 0): Keq increases with temperature. (This is Le Chatelier’s principle in mathematical form.)

Examples:

• H₂(g) + I₂(g) ⇌ 2HI(g): Kc = 794 at 25°C (strongly favors HI)
• N₂(g) + 3H₂(g) ⇌ 2NH₃(g): Kp ≈ 6 × 10⁵ at 25°C (Kc ≈ 3.5 × 10⁸), collapsing to roughly 1.5 × 10⁻⁵ at 500°C. The Haber process runs hot anyway, because at 25°C the equilibrium is superb and the rate is hopeless

What K does and does not tell you

K describes the destination, never the journey. A reaction with a K of 10⁸ ends up almost entirely as products, but it may take a century to get there. Diamond turning into graphite is thermodynamically favourable at room temperature and nobody has ever watched it happen. Rate lives in the activation energy, which K knows nothing about.

The other common misreading is treating K as fixed. It is fixed for a given temperature and nothing else. Changing concentrations or pressures shifts the position of equilibrium, exactly as Le Chatelier describes, but K itself does not move. Change the temperature and K genuinely changes, which is why the Haber numbers above vary so dramatically.

Leaving things out of the expression

Pure solids and pure liquids do not appear in the equilibrium expression. Their concentration is a property of the substance rather than of the mixture, so it is already baked into K. Add more solid calcium carbonate to a decomposition equilibrium and nothing shifts, because you have not changed anything the expression can see. The same goes for water as a solvent in dilute aqueous reactions.

A sanity check on the number

Before trusting a computed K, look at the exponents. Each concentration is raised to its stoichiometric coefficient, so a coefficient of 3 means cubing that term. Get one coefficient wrong and the answer is off by orders of magnitude, which is usually obvious once you compare it against the reference bands.