Acid–base reaction Study Guide

📖 Core Concepts Acid–Base Definitions Arrhenius: Acid → produces H⁺ (hydronium) in water; Base → produces OH⁻. Brønsted–Lowry: Acid = proton donor, Base = proton acceptor; works in any phase. Lewis: Acid = electron‑pair acceptor, Base = electron‑pair donor; emphasizes dative bond formation. Solvent‑System: Acid–base character can flip depending on the solvent (e.g., NH₃ is a base in water but an acid in liquid NH₃). Conjugate Pairs – Removing a proton from an acid gives its conjugate base; adding a proton to a base gives its conjugate acid. Equilibrium Constants – Acid strength: \(Ka\) (larger = stronger acid). Base strength: \(Kb\) (larger = stronger base). Relationship: \(pKa + pKb = 14\) (at 25 °C, aqueous). HSAB Theory – “Hard” species are small & low‑polarizable; “Soft” species are larger & highly polarizable. Like likes like. Thermodynamics – \(\Delta G = -RT\ln K\); \(\Delta G\) determines spontaneity of acid‑base reactions. 📌 Must Remember Arrhenius definitions only apply to aqueous solutions. Water is amphoteric (both acid and base). Strong acid + strong base → reaction goes to completion (only salt + water). Henderson–Hasselbalch: \(\mathrm{pH} = pKa + \log\!\left(\frac{[\text{A}^-]}{[\text{HA}]}\right)\). \(pKa + pKb = 14\) (25 °C, water). HSAB “hard–hard” and “soft–soft” interactions are most stable. \(\Delta G = -RT\ln K\); negative \(\Delta G\) ⇔ \(K > 1\). 🔄 Key Processes Neutralization \(\text{HA}{(aq)} + \text{OH}^-{(aq)} \rightarrow \text{A}^-{(aq)} + \text{H}2\text{O}{(l)}\). Strong acid + strong base → complete ion pairing, spectator ions omitted. Buffer Preparation Choose a weak acid HA and its conjugate base A⁻ with \(pKa\) near desired pH. Adjust ratio using Henderson–Hasselbalch equation. Acid‑Base Titration Add titrant of known concentration until equivalence point (moles acid = moles base). Detect with pH indicator or electrode; calculate unknown concentration: \[ C{\text{unknown}} = \frac{C{\text{titrant}} \times V{\text{titrant}}}{V{\text{unknown}}} \] HSAB Matching Identify hardness/softness of reactants → predict favorability (hard‑hard or soft‑soft). 🔍 Key Comparisons Arrhenius vs. Brønsted–Lowry Arrhenius: limited to water, focuses on H⁺/OH⁻. Brønsted–Lowry: solvent‑independent, focuses on proton transfer. Brønsted–Lowry vs. Lewis Brønsted–Lowry: proton donor/acceptor. Lewis: electron‑pair donor/acceptor (broader; includes reactions without H⁺). Hard vs. Soft Acid/Base Hard: small, high charge, low polarizability (e.g., \( \text{Na}^+ , \text{F}^- \)). Soft: larger, lower charge, high polarizability (e.g., \( \text{I}^- , \text{Pd}^{2+} \)). ⚠️ Common Misunderstandings “All acids donate protons” – true for Brønsted–Lowry, but Lewis acids don’t involve protons. \(pKa + pKb = 14\) always – only valid for conjugate pairs in water at 25 °C. Strong acid = strong base – strength is independent; a strong acid can pair with a weak base (e.g., HCl + NH₃). Buffers resist any pH change – they resist small additions; large amounts overwhelm capacity. 🧠 Mental Models / Intuition “Proton traffic” – Visualize acids as “proton donors” handing off H⁺ to bases (acceptors). “Electron‑pair hand‑shake” – Lewis acids reach out for a lone pair; bases offer it. HSAB “Magnet” – Hard and soft act like magnets that only snap together with the same type. Equilibrium constant ↔ ΔG – Think of \(K\) as a “thermodynamic lever”: \(K > 1\) pulls ΔG negative → reaction favoured. 🚩 Exceptions & Edge Cases Very weak acids/bases (e.g., water) have \(Ka\) or \(Kb\) ≈ \(10^{-14}\); they can act as both acid and base (amphoteric). Non‑aqueous solvents can invert acidity (e.g., HCl is a weak acid in liquid ammonia). Polyprotic acids (H₂SO₄) have two \(Ka\) values; only the first is typically “strong”. Temperature shift – \(pKa\) values change with temperature; the \(pKa + pKb = 14\) rule breaks down outside 25 °C. 📍 When to Use Which Choose Arrhenius when the problem explicitly deals with aqueous H⁺/OH⁻ concentrations. Use Brønsted–Lowry for any proton‑transfer scenario (including gas‑phase or solid‑phase reactions). Apply Lewis when the reaction involves coordination complexes, carbocations, or no explicit H⁺. Employ HSAB to predict stability of metal‑ligand complexes or choose reagents for selective reactions. Henderson–Hasselbalch is the go‑to formula for buffer pH calculations. \(\Delta G = -RT\ln K\) when you need to assess spontaneity from a known equilibrium constant. 👀 Patterns to Recognize Strong acid + strong base → only salt + water (no equilibrium). Weak acid + strong base → buffer region before equivalence; pH rises sharply after. Hard acid + hard base → high‑lattice‑energy salts (e.g., NaF). Lewis acid + Lewis base → formation of a dative covalent bond (e.g., BF₃·NH₃ adduct). pH‑vs‑volume titration curve: flat region (buffer) → steep rise/fall at equivalence → plateau after. 🗂️ Exam Traps Mistaking “base strength” for “basicity” – \(Kb\) reflects equilibrium, not just ability to accept H⁺. Using \(pKa + pKb = 14\) for non‑aqueous or non‑25 °C problems – will give wrong values. Selecting an indicator with pKa far from equivalence pH – leads to premature color change. Assuming all Lewis acids are “strong” – some are weak (e.g., Ag⁺); strength depends on electron‑pair affinity. Confusing amphoteric species – water can appear as both acid (H₂O → H⁺ + OH⁻) and base (H₂O + H⁺ → H₃O⁺); ignoring this may mis‑balance equations. --- This guide condenses the highest‑yield concepts from the outline, giving you a quick‑reference toolkit for any acid–base exam.