PC-06 — Core Concepts in 10 Sentences — Summary

Dynamic equilibrium is the state where the rate of the forward reaction equals the rate of the backward reaction, so concentrations remain constant but the reaction never stops.
The equilibrium constant Kc is the ratio of equilibrium product concentrations to reactant concentrations, each raised to their stoichiometric coefficients, with pure solids and liquids excluded.
For gaseous reactions, Kp relates to Kc by Kp = Kc(RT)^ $\Delta n$ , where $\Delta n$ equals moles of gaseous products minus moles of gaseous reactants.
The reaction quotient Q uses instantaneous concentrations; Q less than K means the reaction proceeds forward, Q greater than K means it proceeds backward.
Le Chatelier's principle states that a system at equilibrium shifts to partially relieve any applied stress, but only temperature changes the value of K itself.
A catalyst lowers activation energy equally for both forward and backward reactions, increasing the rate of attaining equilibrium without shifting the position or changing K.
pH equals −log[ $H^{+}$ ]; for a weak acid, [ $H^{+}$ ] = √(Ka × C); and pH plus pOH equals 14 only at 25°C.
The Henderson-Hasselbalch equation, pH = pKa + log([salt]/[acid]), gives the pH of an acidic buffer and shows that pH equals pKa when salt and acid concentrations are equal.
The common ion effect suppresses ionisation of a weak electrolyte and dramatically reduces the solubility of sparingly soluble salts, with precipitation occurring whenever the ionic product exceeds Ksp.
Three key NEET traps in this topic are: a catalyst never shifts equilibrium, inert gas at constant volume has no effect, and the pH of 10^{-8} M HCl is approximately 6.98, not 8.