Chemical equilibrium
Equilibrium questions test whether you can set up an ICE table, evaluate Kc or Kp from data, predict the direction of shift, and handle heterogeneous systems and solubility products. Approximations must be justified and units handled sensibly.
The equilibrium constant
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 = Kc(RT)^Δn, where Δn counts only gaseous moles. K depends only on temperature; a catalyst changes the rate but never K.
Reaction quotient and direction
Q uses the same expression with current concentrations. If Q < K the net reaction goes forward, if Q > K it goes backward, and Q = K means equilibrium. Use Q, not Le Chatelier alone, when several changes are applied at once.
Le Chatelier's principle
- Concentration: adding a reactant drives the reaction forward.
- Pressure/volume: increasing pressure shifts towards the side with fewer gas moles.
- Temperature: raising T favours the endothermic direction and changes K itself.
- Inert gas: no effect at constant volume; at constant pressure it acts like dilution.
Heterogeneous equilibria and Ksp
For a sparingly soluble salt, Ksp = [cation]^m[anion]^n. The common-ion effect suppresses solubility; compare Q with Ksp to decide whether a precipitate appears. Removing an ion by complexation or acid–base reaction increases solubility.
Worked example
The solubility of AgCl at 25 °C is 1.34 × 10⁻⁵ mol/L. For AgCl(s) ⇌ Ag⁺ + Cl⁻, [Ag⁺] = [Cl⁻] = s, so Ksp = s² = (1.34 × 10⁻⁵)² = 1.8 × 10⁻¹⁰. In 0.010 M NaCl, [Cl⁻] ≈ 0.010 M, so [Ag⁺] = Ksp/0.010 = 1.8 × 10⁻⁸ M - roughly 700 times less soluble, the common-ion effect in action.
Traps
- Putting solids or water into the equilibrium expression.
- Using Kp = Kc without checking Δn or forgetting to convert T to kelvin.
- Claiming a catalyst raises the yield.
- Confusing Ksp with a single concentration rather than a product of concentrations.
- Forgetting that only temperature changes the value of K.