Acid Study Guide

📖 Core Concepts Acid (Brønsted–Lowry) – a species that can donate a proton (H⁺). Lewis Acid – a species that accepts an electron pair to form a covalent bond. Arrhenius Acid – in water, produces hydronium ions (H₃O⁺); shows sour taste, turns blue litmus red, pH < 7. Acid Dissociation: \(\displaystyle \mathrm{HA}\rightleftharpoons \mathrm{H^{+}}+\mathrm{A^{-}}\) Acid Dissociation Constant: \(Ka=\dfrac{[\mathrm{H^{+}}][\mathrm{A^{-}}]}{[\mathrm{HA}]}\) pKₐ: \(pKa=-\log{10}Ka\); lower pKₐ = stronger acid. Conjugate Acid–Base Pair – differ by one proton; HA ↔ A⁻. Strong Acid – dissociates ≈100 % in water (e.g., HCl, HNO₃, H₂SO₄). Weak Acid – only partially dissociates; HA and A⁻ coexist at equilibrium. Polyprotic Acid – possesses > 1 acidic proton (diprotic, triprotic); each step has its own \(K{a1}, K{a2}, …\) with \(K{a1}>K{a2}>K{a3}\). Buffer: mixture of a weak acid and its conjugate base; pH ≈ pKₐ when \([HA]=[A^-]\). Titration Equivalence Point – stoichiometric neutralization of all acidic protons; indicated by a sharp pH change. HSAB Theory – classifies Lewis acids as hard (small, highly charged) or soft (large, polarizable); hardness correlates with acid strength in non‑aqueous media. --- 📌 Must Remember Acid = proton donor (Brønsted) unless “Lewis” is specified. \(Ka\) large ⇔ strong acid; \(pKa\) small ⇔ strong acid. Strong acids: HCl, HBr, HI, HClO₄, HNO₃, H₂SO₄ (first proton). Weak acids dissociate partially; equilibrium constant governs \([H^+]\). pH definition: \(pH = -\log{10}[\mathrm{H^{+}}]\); acidic solution ⇒ \([\mathrm{H^{+}}] > 10^{-7}\ \text{M}\). Polyprotic acids: \(K{a1} > K{a2} > K{a3}\); each successive proton is harder to lose. Buffer region centered at pH = pKₐ of the weak acid. Halogenated carboxylic acids become stronger as more electronegative halogens are attached (F > Cl > H). Superacids are stronger than 100 % H₂SO₄ (pKₐ < −12). HSAB: Hard acids prefer hard bases; soft acids prefer soft bases. --- 🔄 Key Processes Calculate \(Ka\) from pKₐ: \(Ka = 10^{-pKa}\). Find \([H^+]\) for a weak acid: \[ \text{Assume } HA \rightleftharpoons H^+ + A^-;\; Ka = \frac{x^2}{[HA]0 - x}\approx\frac{x^2}{[HA]0} \] Solve for \(x = [H^+]\). Buffer pH (Henderson–Hasselbalch): \(pH = pKa + \log\frac{[A^-]}{[HA]}\). Diprotic acid titration: Add base until first equivalence point (neutralizes \(K{a1}\) proton). Continue to second equivalence point (neutralizes \(K{a2}\) proton). Two buffer regions at \(pH = pK{a1}\) and \(pH = pK{a2}\). Neutralization reaction: Acid + Base → Salt + Water (e.g., \( \mathrm{HCl} + \mathrm{NaOH} \rightarrow \mathrm{NaCl} + \mathrm{H2O}\)). --- 🔍 Key Comparisons Arrhenius vs. Brønsted–Lowry vs. Lewis Arrhenius: acid → H₃O⁺ in water; base → OH⁻. Brønsted–Lowry: acid donates H⁺; base accepts H⁺ (solvent‑independent). Lewis: acid accepts an electron pair; base donates an electron pair. Strong vs. Weak Acid Strong: 100 % dissociation, large \(Ka\), tiny \(pKa\). Weak: partial dissociation, moderate/small \(Ka\), larger \(pKa\). Diprotic vs. Triprotic Acid Diprotic: two \(Ka\) values ( \(K{a1}>K{a2}\) ); two equivalence points. Triprotic: three \(Ka\) values ( \(K{a1}>K{a2}>K{a3}\) ); three equivalence points. Halogenated Carboxylic Acid Strength Fluoro‑ > Chloro‑ > Dichloro‑ > Trichloro‑ (more electron‑withdrawing = stronger acid). Hard vs. Soft Lewis Acid (HSAB) Hard: small, high charge (e.g., \( \mathrm{Al^{3+}} \)). Soft: large, polarizable (e.g., \( \mathrm{Pt^{2+}} \)). --- ⚠️ Common Misunderstandings “Free H⁺ exists in water.” In reality, protons are solvated as \( \mathrm{H3O^+}, \mathrm{H5O2^+}, …\). All strong acids have the same pKₐ. They differ; e.g., HClO₄ (≈ −10) vs. H₂SO₄ (first proton ≈ −3). pKₐ < 0 means the acid is “super”. Superacids are defined relative to 100 % H₂SO₄, not merely negative pKₐ. Buffer pH = pKₐ only when acid is strong. It holds when \([HA]=[A^-]\) regardless of strength. Halogenated acids are all equally strong. Strength scales with number and electronegativity of halogens. All polyprotic acids have equally spaced pKₐ values. They decrease dramatically: \(pK{a1} \ll pK{a2} \ll pK{a3}\). --- 🧠 Mental Models / Intuition “Let‑go‑ability” model: Larger, more electronegative atoms (or resonance‑stabilized conjugate bases) make the H‑A bond weaker → stronger acid. “Buffer midpoint” picture: When exactly half the acid is neutralized, the solution’s pH equals the acid’s pKₐ (visualize a seesaw balanced at the midpoint). “Step‑down ladder” for polyprotic acids: Each rung (Ka) is lower than the one before; think of climbing down a staircase where each step gets smaller. HSAB analogy: Hard acids = “hard‑handed” (grab hard bases); soft acids = “soft‑handed” (prefer soft bases). --- 🚩 Exceptions & Edge Cases Superacids (e.g., fluoroantimonic acid) far exceed the strength of 100 % H₂SO₄. Sulfuric acid: first proton is strong; second proton is only moderate (Ka₂ ≈ 1.2 × 10⁻²). Vinylogous carboxylic acids (e.g., ascorbic acid) are more acidic than regular carboxylic acids due to conjugation across the C=C bond. Polyprotic titration: second equivalence point may be less pronounced if \(K{a2}\) is very weak. --- 📍 When to Use Which Calculate pH of a weak acid: use \(Ka\) expression → quadratic approximation. Design a buffer: choose a weak acid whose pKₐ ≈ target pH; mix with its salt. Choose titration indicator: select one that changes color near the expected equivalence pH (e.g., phenolphthalein for strong acid‑strong base). Predict Lewis acid strength in non‑aqueous media: apply HSAB (hard ↔ hard, soft ↔ soft). Assess carboxylic acid strength: consider halogen substitution (more F/Cl → stronger) or vinylogous structure (increased acidity). --- 👀 Patterns to Recognize \(K{a1} > K{a2} > K{a3}\) for polyprotic acids. pH ≈ pKₐ at the half‑equivalence point of a titration. Sharp pH jump at equivalence for strong acid‑strong base pairs; gradual for weak‑weak pairs. Halogenated carboxylic acids: acidity order follows F > Cl > Br > H and more halogens = stronger acid. HSAB trend: hard acids react rapidly with hard bases (e.g., \( \mathrm{Al^{3+}} + \mathrm{F^-}\)). --- 🗂️ Exam Traps Confusing Arrhenius with Brønsted–Lowry: a question may list “HCl is an Arrhenius acid” – true only in aqueous solution; the broader definition is Brønsted. Assuming all halogenated acids have the same pKₐ: they differ; e.g., chloroacetic (pKₐ ≈ 2.86) vs. trichloroacetic (pKₐ ≈ 0.7). Mix‑up of Ka and pKa signs: larger Ka → smaller pKa; a common distractor flips the relationship. Missing the second equivalence point on a diprotic titration diagram – many students stop at the first jump. Treating vinylogous acids as ordinary carboxylic acids: they are more acidic; answer choices that ignore the double bond are wrong. Superacid definition: selecting any acid with pKₐ < 0 as a superacid; the correct definition requires strength greater than 100 % H₂SO₄ (e.g., fluoroantimonic acid). ---