Chemistry - Reactions and Equilibrium
Chemistry - Reactions and Equilibrium includes stoichiometry, limiting reagents, redox, acid–base calculations, equilibrium constants, and Le Chatelier’s principle. The exam rewards correct mole ratios and clear equilibrium reasoning.
Core rules
- Balance equations before any mole calculation; mass is conserved.
- Moles n = m/M; for gases at the same conditions, volume ratio equals mole ratio.
- Identify the limiting reagent by dividing available moles by the stoichiometric coefficient.
- For aA + bB ⇌ cC + dD, Kc = [C]^c[D]^d/([A]^a[B]^b); pure solids and liquids are omitted.
- A catalyst speeds both directions equally and does not shift equilibrium.
Worked example
For H₂(g) + I₂(g) ⇌ 2HI(g), a mixture at equilibrium has [H₂] = 0.20 mol dm⁻³, [I₂] = 0.20 mol dm⁻³, and [HI] = 1.6 mol dm⁻³. Calculate Kc. Kc = [HI]²/([H₂][I₂]) = (1.6)²/(0.20 × 0.20) = 2.56/0.040 = 64. The exponent 2 comes from the coefficient of HI. If the reaction were reversed, Kc would be the reciprocal, 1/64.
Traps
- Using masses directly in Kc instead of equilibrium concentrations.
- Including a solid such as CaCO₃ in the equilibrium expression.
- Saying a catalyst increases yield; it only reaches equilibrium faster.
- Forgetting that increasing pressure shifts equilibrium toward the side with fewer gas moles, if those differ.
- Using initial concentrations unchanged in an ICE table.