Acid-Base Theories: Bronsted-Lowry & Lewis JEE Guide

JEE Main tests acid-base chemistry across three different theoretical frameworks — Arrhenius, Bronsted-Lowry, and Lewis — and requires you to apply them in both conceptual and numerical contexts. The Ionic Equilibrium chapter contributes 3–4 questions per session, with acid-base theory, pH calculations, buffer chemistry, and hydrolysis of salts being the core subtopics. This guide covers the conceptual foundation systematically before the numerical applications.

Arrhenius, Bronsted-Lowry, and Lewis: The Three Frameworks

Arrhenius acid: produces H⁺ in water. Arrhenius base: produces OH⁻ in water. Limited to aqueous systems. Bronsted-Lowry acid: proton (H⁺) donor. Bronsted-Lowry base: proton acceptor. Conjugate acid-base pairs: differ by one H⁺. NH₃ + H₂O ⇌ NH₄⁺ + OH⁻; conjugate pairs are NH₃/NH₄⁺ and H₂O/OH⁻. Every Bronsted-Lowry acid reaction has an acid on one side and its conjugate base on the other — identify both in every equilibrium. Strong acid has a weak conjugate base (HCl/Cl⁻) and vice versa (weak acid has a strong conjugate base, CH₃COOH/CH₃COO⁻). This inverse relationship underlies buffer chemistry and salt hydrolysis.

Lewis acid: electron pair acceptor. Lewis base: electron pair donor. This framework includes reactions with no proton transfer. Examples: BF₃ + NH₃ → F₃B←NH₃ (Lewis acid-base adduct). H⁺ is a Lewis acid (accepts electron pair). Metal cations are Lewis acids (form complex ions with Lewis base ligands). AlCl₃, BF₃, SO₃, and carbocations are common Lewis acids in JEE questions. H₂O, NH₃, amines, and all anions with lone pairs are Lewis bases. Test your acid-base theory with a free mock before the numerical section. For the full ionic equilibrium framework, see our ionic equilibrium guide.

Relative Acid and Base Strength

Factors affecting acid strength: (1) bond polarity and bond length — HF > HCl for electronegativity within a row, but HI > HBr > HCl for bond strength/enthalpy down a group; (2) electron-withdrawing groups stabilise the conjugate base → increase acid strength (trifluoroacetic acid >> acetic acid); (3) resonance stabilisation of conjugate base — carboxylic acids are stronger than alcohols because the carboxylate anion is resonance-stabilised; (4) inductive effect and field effect. For oxyacids of the same element: more oxygen atoms → more acidic (H₂SO₄ > H₂SO₃; HClO₄ > HClO₃ > HClO₂ > HClO). For oxyacids with same oxidation state different central element: same group → lower row more acidic (H₂SO₄ > H₂SeO₄).

Base strength factors: lone pair availability, steric hindrance, resonance. Aliphatic amines are stronger bases than aniline (lone pair of aniline is delocalised into ring). Ethylamine > methylamine in terms of proton donation ease (inductive effect) for gas phase; in water, solvation effects complicate the order and make secondary amines the strongest for small alkyl groups (methylamine series: dimethylamine > methylamine > trimethylamine in water). For amine chemistry, see our nitrogen compounds guide.

pH Calculations: The Core Numericals

Strong acid (HA) of molarity C: [H⁺] = C, pH = −log C. For very dilute strong acid (C < 10⁻⁶ M), include water autoionisation: [H⁺] = C + 10⁻⁷ + correction (solve quadratic). Weak acid (Ka): [H⁺] = √(Ka × C) when α << 1 (degree of dissociation small); pH = ½(pKa − log C). Strong base (BOH) of molarity C: [OH⁻] = C, pOH = −log C, pH = 14 − pOH. Buffer (Henderson-Hasselbalch): pH = pKa + log([A⁻]/[HA]). Buffer capacity is maximum at pH = pKa (equal concentrations of acid and conjugate base). These five formulae cover 90% of all pH JEE questions.

Hydrolysis of salts: salt of strong acid + weak base (NH₄Cl) → acidic solution; pH = 7 − ½(pKb − log C). Salt of weak acid + strong base (CH₃COONa) → basic solution; pH = 7 + ½(pKa + log C). Salt of weak acid + weak base: pH = 7 + ½(pKa − pKb). These three hydrolysis formulae are frequently tested. For complete acid-base equilibrium coverage, see our buffer solutions and pH guide.

Amphoteric Species and Polyprotic Acids

Amphoteric species act as both acid and base: H₂O, HCO₃⁻, HSO₄⁻, H₂PO₄⁻, HPO₄²⁻, HS⁻, amino acids (zwitterionic form). Isoelectric point of an amino acid: pH at which net charge is zero — exactly between pKa values for simple amino acids. Polyprotic acids (H₃PO₄, H₂SO₄, H₂CO₃) ionise stepwise with Ka₁ >> Ka₂ >> Ka₃. For polyprotic weak acid, [H⁺] ≈ √(Ka₁ × C). The intermediate ion of a polyprotic acid: pH = ½(pKa₁ + pKa₂). These results are tested in about two JEE questions per year and require careful application. For the broader equilibrium framework, see our chemical equilibrium guide.