Rucete ✏ AP Chemistry In a Nutshell
13. Acids and Bases
This chapter explains acid–base theories, strength comparisons, pH and pOH calculations, hydrolysis, salt solutions, buffers, titration curves, and acid–base equilibria using Ka, Kb, and the Henderson–Hasselbalch equation.
Acid–Base Theories
• Arrhenius: Acids increase H⁺; bases increase OH⁻ in water.
• Brønsted–Lowry: Acids donate protons (H⁺); bases accept protons.
• Acids have weakly bound hydrogens that ionize in water.
Acid and Base Nomenclature
• Binary acids: hydro- prefix and -ic suffix (e.g., HCl → hydrochloric acid).
• Polyatomic acids: – "-ate" becomes "-ic" (e.g., HNO₃ → nitric acid) – "-ite" becomes "-ous" (e.g., HNO₂ → nitrous acid)
• Organic acids: names end in -oic acid (e.g., CH₃COOH → ethanoic acid).
• Bases with OH⁻: named as "metal + hydroxide".
• Nitrogen-containing bases are amines (e.g., CH₃NH₂ → methylamine).
Strong and Weak Acids
• Strong acids fully ionize in water: HCl, HBr, HI, HNO₃, HClO₄, H₂SO₄ (first H⁺ only).
• Weak acids only partially ionize: HF, H₃PO₄, H₂CO₃, CH₃COOH, etc.
• Acid strength increases: – across a period (due to electronegativity) – down a group (due to increasing anion size and weaker H–X bonds)
• In oxoacids, more oxygen atoms and more electronegative central atoms → stronger acids.
Strong and Weak Bases
• Strong bases: Group 1 and some Group 2 metal hydroxides (e.g., NaOH, Ba(OH)₂).
• Weak bases: ammonia (NH₃) and amines (e.g., CH₃NH₂, C₂H₅NH₂).
• Electronegative groups reduce base strength in organic amines.
Acidic and Basic Anhydrides
• Acid anhydrides: nonmetal oxides → acids when added to water (e.g., SO₃ + H₂O → H₂SO₄).
• Basic anhydrides: metal oxides → bases when added to water (e.g., CaO + H₂O → Ca(OH)₂).
Neutralization Reactions
• Acid + base → salt + water (e.g., HCl + NaOH → NaCl + H₂O).
• Reaction can be written as molecular, complete ionic, or net ionic equations.
• Polyprotic acid neutralization occurs stepwise depending on OH⁻ ratio.
Polyprotic Acids
• Acids with multiple ionizable H⁺ (e.g., H₂SO₄, H₃PO₄).
• Ionization occurs in steps; each step has a separate Ka.
• Neutralization with base forms stepwise phosphate salts depending on mole ratio.
Conjugate Acid–Base Pairs
• Differ by one proton (e.g., HCl / Cl⁻, NH₄⁺ / NH₃).
• Strength relationship: – Strong acid → very weak conjugate base – Weak acid → stronger conjugate base
• Water is amphiprotic: can act as acid or base (H₂O ⇌ H⁺ + OH⁻).
Autoionization of Water
• Water self-ionizes slightly: H₂O ⇌ H⁺ + OH⁻
• At 25°C, Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴
• In pure water, [H⁺] = [OH⁻] = 1.0 × 10⁻⁷ M
pH and pOH
• pH = –log[H⁺], pOH = –log[OH⁻]
• pH + pOH = 14.00 at 25°C
• [H⁺] = 10⁻ᵖᴴ, [OH⁻] = 10⁻ᵖᴼᴴ
• Acidic: pH < 7, Neutral: pH = 7, Basic: pH > 7
Ka and Kb
• Ka = acid ionization constant: HA ⇌ H⁺ + A⁻
• Kb = base ionization constant: B + H₂O ⇌ BH⁺ + OH⁻
• Strong acids/bases: no Ka/Kb used since they fully dissociate
• Weak acids/bases: Ka and Kb used to find pH or concentrations
Ka and Kb Relationship
• Ka × Kb = Kw = 1.0 × 10⁻¹⁴
• pKa + pKb = 14.00
• Knowing Ka lets you find Kb for its conjugate base (and vice versa)
Salts and pH
• Neutral salts: strong acid + strong base (e.g., NaCl)
• Acidic salts: strong acid + weak base (e.g., NH₄Cl)
• Basic salts: weak acid + strong base (e.g., CH₃COONa)
• Some salts hydrolyze to change pH; calculate using Ka or Kb of the conjugate ion
Buffer Solutions
• Made of weak acid and its salt or weak base and its salt
• Resist pH change when small amounts of acid or base are added
• Common examples: – CH₃COOH / CH₃COONa – NH₃ / NH₄Cl
Henderson–Hasselbalch Equation
• For buffers: pH = pKa + log([A⁻]/[HA])
• For bases: pOH = pKb + log([BH⁺]/[B])
• Best buffering when [A⁻] ≈ [HA] and pH ≈ pKa
Titration Curves
• pH vs. volume of titrant added
• Strong acid + strong base → sharp equivalence point at pH 7
• Weak acid + strong base → equivalence point > 7
• Strong acid + weak base → equivalence point < 7
• Half-equivalence point: pH = pKa
Indicators
• Weak acids or bases with different colors in protonated/deprotonated forms
• Change color near pKa value
• Choose indicator with pKa ≈ equivalence point pH
In a Nutshell
Acids and bases are defined by their ability to donate or accept protons, and their strengths affect solution pH. Weak acids and bases form equilibria that can be quantified with Ka and Kb, while buffers and titration curves help maintain and analyze pH in reactions. Understanding conjugate pairs, salt hydrolysis, and indicators is key to mastering acid–base chemistry.