Reactivity 3.1 Flashcards

A proton (H⁺) donor.

A proton (H⁺) acceptor.

Two species that differ by exactly one proton. E.g. HCl/Cl⁻ or NH₄⁺/NH₃.

Pair 1: NH₃ (base) / NH₄⁺ (conjugate acid). Pair 2: H₂O (acid) / OH⁻ (conjugate base).

HSO₄⁻ (remove one H⁺ and decrease charge by 1).

A substance that can both donate and accept a proton (H⁺). Examples: H₂O, HCO₃⁻, HSO₄⁻, amino acids.

Amphiprotic = can donate/accept H⁺ (Brønsted-Lowry). Amphoteric = can react as acid or base (broader, includes Lewis). All amphiprotic species are amphoteric, but not vice versa (e.g. Al₂O₃ is amphoteric but not amphiprotic).

As acid: H₂O + NH₃ ⇌ NH₄⁺ + OH⁻ (donates H⁺). As base: H₂O + HCl → H₃O⁺ + Cl⁻ (accepts H⁺).

Salt + water. E.g. HCl + NaOH → NaCl + H₂O.

Salt + water + CO₂. E.g. 2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂.

Strong acids fully dissociate (→) — nearly 100% ionised. Weak acids partially dissociate (⇌) — typically <5% ionised. Strength ≠ concentration.

HCl, HBr, HI, HNO₃, H₂SO₄ (1st proton only).

Group 1 hydroxides (LiOH, NaOH, KOH, RbOH, CsOH) + heavy Group 2 hydroxides (Ca(OH)₂, Sr(OH)₂, Ba(OH)₂).

Strength = extent of ionisation (all or partial). Concentration = moles of solute per dm³. A concentrated weak acid (e.g. 1.0 M CH₃COOH, pH ≈ 2.4) can have a lower pH than a dilute strong acid (e.g. 0.0001 M HCl, pH = 4.0), showing that pH depends on both concentration and strength.

pH = −log₁₀[H⁺]. Therefore [H⁺] = 10^⁻ᵖᴴ.

pOH = −log₁₀[OH⁻]. At 25°C: pH + pOH = 14.00.

The ionic product of water: Kw = [H⁺][OH⁻] = 1.00 × 10⁻¹⁴ at 25°C. In pure water: [H⁺] = [OH⁻] = 1.0 × 10⁻⁷ M → pH = 7.

HNO₃ is a strong acid → fully dissociates. [H⁺] = 0.050 M. PH = −log(0.050) = 1.30.

NaOH is a strong base → [OH⁻] = 0.020 M. POH = −log(0.020) = 1.70. PH = 14.00 − 1.70 = 12.30.

pH = 7, because the salt formed (e.g. NaCl) does not hydrolyse.

Auto-ionisation of water is endothermic (ΔH > 0). Higher T shifts equilibrium right → more H⁺ and OH⁻ → Kw increases → pH of pure water decreases (but water is still neutral because [H⁺] = [OH⁻]).

Still neutral. [H⁺] = √(2.4 × 10⁻¹⁴) = 1.55 × 10⁻⁷ M. PH = 6.81. Neutral ≠ pH 7; neutral = [H⁺] = [OH⁻].

Ka = acid dissociation constant = [H⁺][A⁻]/[HA]. PKa = −log Ka. Larger Ka (smaller pKa) = stronger weak acid.

Kb = base dissociation constant = [BH⁺][OH⁻]/[B]. PKb = −log Kb. Larger Kb (smaller pKb) = stronger weak base.

Ka × Kb = Kw. Therefore pKa + pKb = pKw = 14.00 (at 25°C). Stronger acid → weaker conjugate base.

Ka = x²/0.10 → x = √(1.8 × 10⁻⁶) = 1.34 × 10⁻³ M. PH = −log(1.34 × 10⁻³) = 2.87.

pH > 7. The conjugate base (e.g. CH₃COO⁻) hydrolyses in water to produce OH⁻, making the solution basic.

pH < 7. The conjugate acid (e.g. NH₄⁺) hydrolyses to produce H⁺.

Neither the acid nor base fully ionises, so there is no single dramatic pH jump. A pH meter is needed; no indicator works.

The indicator’s pKa (colour-change range) must fall within the steep vertical section of the titration curve. Phenolphthalein (8.3–10.0) for WA+SB. Methyl orange (3.1–4.4) for SA+WB.

When dissolved salt ions react with water, affecting pH. SA+SB salt → pH 7. WA+SB salt → pH > 7. SA+WB salt → pH < 7.

A solution that resists changes in pH when small amounts of acid or base are added. Made from comparable amounts of a weak acid + its conjugate base (acidic buffer) or weak base + its conjugate acid (basic buffer).

If H⁺ added: A⁻ + H⁺ → HA (conjugate base absorbs it). If OH⁻ added: HA + OH⁻ → A⁻ + H₂O (weak acid neutralises it).

pH = pKa + log([A⁻]/[HA]). At the half-equivalence point [A⁻] = [HA], so pH = pKa.

At half-equivalence, exactly half the acid has been neutralised, so [HA] = [A⁻]. Log([A⁻]/[HA]) = log(1) = 0. Therefore pH = pKa + 0 = pKa.