How reversible reactions balance and how pH is calculated
In a reversible reaction in a closed system, the forward and backward reactions continue at equal rates once equilibrium is reached, so concentrations stop changing even though the reactions have not stopped. This is dynamic equilibrium. For aA + bB ⇌ cC + dD, the equilibrium constant is Kc = [C]ᶜ[D]ᵈ/[A]ᵃ[B]ᵇ, and for gases Kp = Kc(RT)^Δn. Le Chatelier’s principle predicts that a system shifts to oppose any change in concentration, pressure or temperature.
Ionic equilibrium applies these ideas to acids and bases. Water ionises slightly, with Kw = [H⁺][OH⁻] = 1 × 10⁻¹⁴ at 298 K, and pH = −log[H⁺]. Weak acids have dissociation constants Ka, and the degree of ionisation follows Ostwald’s dilution law. Buffer solutions resist pH change; for an acidic buffer, pH = pKa + log([salt]/[acid]). The common ion effect suppresses ionisation, and a precipitate forms when the ionic product exceeds the solubility product Ksp.
Why it matters
The Haber process for ammonia is run at conditions chosen using Le Chatelier’s principle. Blood stays near pH 7.4 because of natural buffers, and medicines, shampoos and soil treatments are all pH-controlled. Solubility product explains kidney stone formation and helps chemists separate ions in analysis.
Key terms
- Dynamic equilibrium
- A state where forward and reverse reactions occur at equal rates, so concentrations stay constant.
- Le Chatelier’s principle
- A system at equilibrium shifts to counteract any change imposed on it.
- Buffer solution
- A solution that resists changes in pH when small amounts of acid or base are added.
- Solubility product
- The product of ion concentrations in a saturated solution of a sparingly soluble salt.