Chapter 8: Acids and Bases

Acid: Proton (H+) donor Base: Proton (H+) acceptor

Conjugate acid-base pairs: Substances that differ by one proton (H+) Conjugate bases: Formed when acids donate a proton (H+), allowing it to act as a base in the reverse reaction Conjugate acids: Formed when bases accept a proton (H+), allowing it to act as an acid in the reverse reaction

Amphiprotic: Substance that can both accept and donate a proton, and can hence act as both a base and an acid Amphoteric: Substance that can react with both acids and bases, not necessarily involving proton transfer (all amphiprotic substances are amphoteric, but not all amphoteric substances are amphiprotic)

Salt: Neutral substance made up of a metal cation (from the base) and non-metal anion (from the acid)

If base is in excess, the diprotic acid is able to undergo two-step dissociation: Eg. H₂SO₄ + 2 NaOH –> Na₂SO₄+ 2H₂O If base is limiting, the diprotic acid can only undergo its first dissociation: Eg. H₂SO₄+ 1 NaOH –> NaHSO₄+ H₂O

• Acids + Reactive metal –> Salt + H₂ (g) (*Link to Redox Processes: Reactive metals are metals higher than H on the activity series, reducing H+ to H₂) • Acids + Carbonates –> Salt + H₂O (l) + CO₂ (g) • Acids + Bases –> Salt + H₂O (l) [Neutralisation –> exothermic reaction]

Acidic: pH<7 at 25°C Neutral: pH=7 at 25°C Alkaline: pH>7 at 25°C

pH: simplified expression of the concentration of H+ ions in a solution (not in mol dm⁻³) pH = - log₁₀[H⁺] [H⁺] = 10⁻ᵖᴴ

pH 0 = 1mol dm⁻³ [H⁺] pH 1 = 0.1 mol dm⁻³ [H⁺] pH 2 = 0.01 mol dm⁻³ [H⁺] 0.2 mol dm⁻³ [H⁺] ≠ pH 2

[H⁺] in 10cm³ = 0.1 mol dm⁻³ + 10cm³ of water –> 2x dilution [H⁺] in 20cm³ = 0.05 mol dm⁻³ pH will increase

H₂O (l) ⇌ H⁺ (aq) + OH⁻ (aq)

Kᵥᵥ = 1.0 x 10⁻¹⁴ [H⁺] [OH⁻] = 1.0 x 10⁻¹⁴ [H⁺] = 1.0 x 10⁻¹⁴/[OH⁻] [OH⁻] = 1.0 x 10⁻¹⁴/[H⁺] Acidic: [H⁺] > [OH⁻] Neutral: [H⁺] = [OH⁻] Basic: [H⁺] < [OH⁻]

Since HCI is a strong acid, assume complete dissociation: [HCI] = [H⁺] = 0.10 mol dm⁻³ Since [H⁺] [OH⁻] = 1.0 x 10⁻¹⁴, [OH⁻] = 1.0 x 10⁻¹⁴/0.10 = 1.0 x 10⁻¹³ Hence, HCI is acidic, as [H⁺] > [OH⁻]

Acids: Dissociate in water by donating a proton (H⁺) to water HA (aq) + H₂O (l) ⇌ A (aq) + H₃O (aq) Strong acids: Dissociate completely in aqueous solution (equilibrium lies long way to the right, allowing us to assume [HA] = [H⁺] Eg. HCI, H₂SO₄, HNO₃ Weak acids: Dissociate partially in aqueous solution (equilibrium lies to the left: [HA] ≠ [H⁺] ) Eg. H₂CO₃, CH₃COOH Bases: Ionise in water by receiving a proton (H⁺) from water B (aq) + H₂O (l) ⇌ BH⁺ (aq) + OH⁻ (aq) Strong base: Ionise completely in aqueous solution; [B]=[OH⁻] Eg. Group 1 metal hydroxides and Ba(OH)₂ Weak base: Ionise partially in aqueous solution [B] ≠ [OH⁻] Eg. NH₃

• Strength: Degree of dissociation • Proticity: Number of H atoms that can be donated Eg. - Both HCI (monoprotic) and H₂SO₄ (diprotic) are strong acids - H₂SO₄ and H₂CO₃ are both diprotic, but H₂SO₄ is strong, and H₂CO₃ is weak

The stronger the acid/base, the weaker its conjugate base/acid - Strong acids/bases: Dissociate/ionise completely in aqueous solution –> equilibrium lies long way to the right –> conjugate base/acid unlikely to reform parent acid/base The weaker the acid/base, the stronger its conjugate base/acid - Weak acids/bases: Dissociate/ionise partially in aqueous solution –> equilibrium lies to the left –> high tendency for conjugate base/acid to reform parent acid/base

*Only acids/bases of equal concentrations can be compared (concentrated weak acid can have a higher [H⁺] than a diluted strong acid): • Strong acids are better conductors of electricity than weak acids (dissociates completely in aqueous solution = greater concentration of ions = more charge carriers): Conductivity meter/lightbulb in an electrolytic cell set-up • Strong acids will have a lower pH than weak acids (dissociate completely = higher [H⁺] = higher pH): measure using pH probe/universal indicator • Strong acids will react faster and more vigorously with metals to produce H₂: More rapid bubbling will be observed

Strength: Degree of dissociation Concentration: Number of moles of acid before dissociation in mol dm⁻³

Equal concentrations of acids, regardless of their strength, will react with equal number of moles of base

H₂O (l) + CO₂ (g) –> H₂CO₃ (aq) [carbonic acid: weak acid] pH range: 5.6 to 7

Acid deposition: when acidic substances leave the earth’s atmosphere and are deposited onto the earth’s surface (wet and dry deposition) Acidic pollutants: sulfur and nitrogen oxides (from coal power plants and combustion engines) S (s) + O₂ (g) –> SO₂ (aq) 2 SO₂ (aq) + O₂ –> SO₃ (aq) SO₃ (aq) + H₂O (l) –> H₂SO₄ (aq) [sulfuric acid] N₂ (g) + O₂ (g) –> 2 NO (g) 2 NO (g) + O₂ (g) –> 2 NO₂ (g) 2 NO₂ (g) + H₂O (l) –> HNO₂ + HNO₃ [nitric acid]

Vegetation: H⁺ ions in acid rain displace the metal ions in soil, which are crucial for the survival of vegetation (Eg. Mg²⁺ ions are crucial for plants to produce chlorophyll, causing them to be unable to photosynthesise properly if displaced) Aquatic life: cannot survive below certain pH levels Marble/limestone structures: Eroded by acid rain (acid-carbonate reaction) CaCO₃ (s) + H₂SO₄ –> CaSO₄ + H₂O + CO₂ Human health: mucosal membranes are irritated by acids, causing respiratory illnesses. acid rain can also react with heavy metal compounds and release poisonous ions

Pre-combustion desulfurisation: Remove sulfur before fuel is burned - Heat crude oil fractions with hydrogen in the presence of a catalyst, to produce H₂S, which can be removed by bubbling gas through an alkaline or converting it back to S, which can be sold to companies to manufacture sulfuric acid Post-combustion desulfurisation: Remove sulfur after fuel is burned - Pass gases through a vessel, where SO₂ can react with calcium oxide/carbonate/hydroxide - Liming of lakes (using Ca(OH)₂ )