Master one of chemistry's most important topics with this comprehensive acids, bases, and pH flashcard deck. Understanding acid-base chemistry is essential for biology, medicine, environmental science, and industrial chemistry — and it is one of the most heavi...
Master one of chemistry's most important topics with this comprehensive acids, bases, and pH flashcard deck. Understanding acid-base chemistry is essential for biology, medicine, environmental science, and industrial chemistry — and it is one of the most heavily tested topics in AP Chemistry, A-Level, GCSE, and university general chemistry courses.
This deck covers the three major definitions of acids and bases (Arrhenius, Brønsted-Lowry, and Lewis), the pH scale, strong and weak acids and bases, neutralization reactions, buffer solutions, titrations, Ka and Kb values, and common acid-base indicators. Includes worked examples of pH calculations.
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Acid: Produces H⁺ ions in water. Base: Produces OH⁻ ions in water. Example: HCl → H⁺ + Cl⁻. NaOH → Na⁺ + OH⁻. Simplest definition but limited to aqueous solutions.
Acid: A proton (H⁺) donor. Base: A proton (H⁺) acceptor. More general than Arrhenius — applies to non-aqueous solutions too. Example: HCl donates H⁺ to NH₃ (base).
Lewis acid: An electron pair acceptor. Lewis base: An electron pair donor. Broadest definition — explains reactions with no proton transfer. Example: BF₃ (Lewis acid) + NH₃ (Lewis base).
pH measures the concentration of H⁺ ions in solution. Scale runs from 0–14. pH < 7 = acidic. pH = 7 = neutral. pH > 7 = basic/alkaline. Each unit is a 10-fold change in H⁺ concentration.
pH = −log[H⁺] (where [H⁺] is the molar concentration of hydrogen ions). Example: [H⁺] = 0.01 mol/L → pH = −log(0.01) = 2.
pOH = −log[OH⁻]. At 25°C: pH + pOH = 14. Example: If pH = 3, then pOH = 11.
Strong acids completely dissociate in water. The 6 common strong acids: HCl (hydrochloric), HBr (hydrobromic), HI (hydroiodic), HNO₃ (nitric), H₂SO₄ (sulfuric), HClO₄ (perchloric).
Weak acids partially dissociate in water — establishing an equilibrium. Examples: CH₃COOH (acetic/ethanoic acid), H₂CO₃ (carbonic acid), HF (hydrofluoric acid), H₃PO₄ (phosphoric acid).
Weak bases partially ionize in water. Examples: NH₃ (ammonia), CH₃NH₂ (methylamine), Na₂CO₃ (sodium carbonate). Form equilibrium with their conjugate acid.
An acid and base that differ by a single proton (H⁺). When an acid donates H⁺, it becomes its conjugate base. Example: HCl/Cl⁻ and H₂O/OH⁻ are conjugate pairs.
The reaction between an acid and a base to form salt + water. Example: HCl + NaOH → NaCl + H₂O. In all neutralizations: H⁺ + OH⁻ → H₂O.
A measure of the strength of a weak acid. Larger Ka = stronger acid (more dissociated). Ka = [H⁺][A⁻] / [HA]. For strong acids, Ka is very large (essentially complete dissociation).
A solution that resists changes in pH when small amounts of acid or base are added. Contains a weak acid and its conjugate base (or weak base and conjugate acid). Example: CH₃COOH / CH₃COO⁻.
pH = pKa + log([A⁻]/[HA]). Used to calculate the pH of a buffer solution. pH equals pKa when [A⁻] = [HA] (equal concentrations of acid and conjugate base).
A technique to determine the concentration of an unknown acid or base by slowly adding a solution of known concentration (titrant) until the reaction is complete (equivalence point).
The point at which the moles of acid exactly equal the moles of base — the reaction is complete. Detected using an indicator or pH meter.
The point in a weak acid titration where exactly half the acid has been neutralized. At this point, pH = pKa — the buffer capacity is greatest.
Litmus: Red in acid, blue in base. Phenolphthalein: Colorless in acid, pink/red in base (changes at pH 8.2–10). Methyl orange: Red in acid, yellow in base (changes at pH 3.1–4.4).
Water can act as both an acid and a base: H₂O ⇌ H⁺ + OH⁻. At 25°C, Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴. In pure water, [H⁺] = [OH⁻] = 1.0 × 10⁻⁷ mol/L → pH = 7.
A substance that can act as both an acid and a base. Water is the most common example. Amino acids and bicarbonate (HCO₃⁻) are also amphoteric.
Rain with pH < 5.6 caused by SO₂ and NO₂ (from burning fossil fuels) dissolving in rainwater to form H₂SO₄ (sulfuric acid) and HNO₃ (nitric acid). Damages ecosystems and stone structures.
Dilute: Low concentration of acid in water (e.g., 0.1 M HCl). Concentrated: High concentration of acid (e.g., 12 M HCl). NOTE: Dilute vs concentrated is independent of strong vs weak.
The reaction of salt ions with water to produce an acidic or basic solution. Salts from strong acid + weak base = acidic solution. Weak acid + strong base = basic solution. Strong acid + strong base = neutral.
pKa = −log(Ka). A lower pKa means a stronger acid (more dissociated). Example: HCl pKa ≈ −7 (very strong). Acetic acid pKa = 4.76 (weak acid).
