Equilibrium
Class 11 · Updated 2026-10-03. Published for practice.
Dynamic equilibrium
At equilibrium, forward and reverse reactions continue but net concentrations stay constant. Macroscopically nothing changes; microscopically molecules still react both ways.
Equilibrium can be reached from either direction. Only closed systems at constant temperature reach a stable equilibrium composition.
A catalyst speeds both directions equally. It helps you reach equilibrium faster; it does not change the equilibrium position or K.
Law of mass action
For a general reaction aA + bB ⇌ cC + dD, Kc = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ at equilibrium, with concentrations in mol L-1. Pure solids and liquids are omitted from the expression.
Kp uses partial pressures for gases: Kp = Kc(RT)^Δn when Δn = (moles gaseous products − moles gaseous reactants) from the balanced equation.
K is fixed at a given temperature. If Q (reaction quotient with current concentrations) < K, the reaction shifts forward; if Q > K, it shifts reverse until Q = K.
Le Chatelier’s principle
If a stress is applied to a system at equilibrium, the system shifts to partially oppose that stress. Stresses you meet in Class 11: concentration change, pressure change (gases), temperature change.
Adding reactant shifts toward products; removing product does the same. For gases, increasing pressure favours the side with fewer moles of gas (if Δn ≠ 0).
Temperature: treat heat as a product in endothermic reactions and a reactant in exothermic ones. Heating favours the endothermic direction. K itself changes with T.
Ionic equilibrium (intro)
Weak acids and bases do not fully ionise. Ka and Kb measure that partial ionisation. Kw = [H+][OH-] = 10-14 at 25 °C links acid and base ionisation in water.
pH = −log10[H+] (for dilute aqueous solutions where activity ≈ concentration). Strong acid/base numericals are stoichiometry; weak acid pH needs Ka and often the approximation [H+] ≈ √(Ka × C) when ionisation is small.
Common ion effect: adding a salt that shares an ion with a weak electrolyte suppresses ionisation (e.g. NaCH3COO added to CH3COOH lowers [H+]). Buffer action builds on this idea in higher classes.