Equilibrium
Equilibrium: constants, quotients, and shifts
At equilibrium, forward and reverse rates are equal and concentrations are constant-but not equal. For aA + bB ⇌ cC + dD, Kc = [C]^c[D]^d/([A]^a[B]^b). Pure solids and liquids are omitted because their activity is 1. Kp uses partial pressures; Kp = Kc(RT)^Δn, where Δn is moles of gaseous products minus gaseous reactants.
The reaction quotient Q has the same form as K but uses current concentrations. If Q < K, the reaction proceeds right; if Q > K, it proceeds left; if Q = K, it is at equilibrium. Manipulating equations: reversing inverts K, multiplying coefficients raises K to that power, and adding reactions multiplies their K values.
Le Châtelier's principle: adding reactant shifts right; removing product shifts right; decreasing volume shifts toward the side with fewer gas moles; changing temperature changes K itself (endothermic reactions favor products at higher T). A catalyst speeds approach to equilibrium but changes neither K nor the equilibrium position. Adding an inert gas at constant volume has no effect.
Traps: pressure changes only matter when gas mole counts differ; solids do not appear in Q or K; and a small K does not mean the reaction is slow-it means products are poorly favored.
Worked example: Kc = 4.0 for A ⇌ B. If [A] = 2.0 M and [B] = 1.0 M, Q = 0.50 < K, so the reaction shifts right.
Rule of thumb: compute Q first, then compare it with K to predict direction.