Chapter 6: Acid, Base and Salt

6.1 The Role of Water in Showing Acidic and Alkaline Properties

• Acid (Arrhenius definition): a substance that ionises in water to produce hydrogen ions, H⁺. (p.136)
• Basicity of an acid = number of H⁺ ions produced per molecule: monoprotic (e.g. HCl), diprotic (e.g. H₂SO₄), triprotic (e.g. H₃PO₄). Base reacts with acid to form salt and water only; an alkali is a base that is soluble in water and ionises to produce hydroxide ions, OH⁻. (p.137)
• Acids and alkalis only show their acidic/alkaline properties when dissolved in water (water enables ionisation to release H⁺ or OH⁻). (p.140)

6.2 pH Value

• pH scale ranges from 0 to 14; pH < 7 is acidic, pH > 7 is alkaline, pH = 7 is neutral. pH = −log[H⁺]. (p.143)
• pOH = −log[OH⁻], and pH + pOH = 14. The higher the concentration of H⁺, the lower the pH; the higher the concentration of OH⁻, the higher the pH. (p.144)

6.3 Strength of Acids and Alkalis

• Strong acid/alkali ionises completely in water to produce a high concentration of H⁺/OH⁻ (e.g. HCl, NaOH). Weak acid/alkali ionises only partially, producing a low concentration of H⁺/OH⁻ (e.g. CH₃COOH, NH₃). (p.150)

6.4 Chemical Properties of Acids and Alkalis

• Acid + Base → Salt + Water; Acid + Reactive metal → Salt + Hydrogen gas; Acid + Metal carbonate → Salt + Water + Carbon dioxide gas. (p.157)
• Alkali + Acid → Salt + Water; Alkali + Ammonium salt → Salt + Water + Ammonia gas (on heating); Alkali + Metal ion → Insoluble metal hydroxide precipitate + Cation from the alkali. (p.157)

6.5 Concentration of Aqueous Solution

• Concentration (g dm⁻³) = mass of solute (g) / volume of solution (dm³). Molarity (mol dm⁻³) = number of moles of solute / volume of solution (dm³). (p.158)

6.6 Standard Solution

• A standard solution is a solution of known (accurately determined) concentration, prepared either by dissolving a weighed mass of solid or by diluting a stock solution. (p.162)
• Dilution formula: M₁V₁ = M₂V₂ (number of moles of solute is unchanged by dilution, only concentration changes). (p.164)

6.7 Neutralisation

• Neutralisation is the reaction between an acid and an alkali to produce salt and water only; ionic equation: H⁺(aq) + OH⁻(aq) → H₂O(l). (p.168)
• Titration is a quantitative method using an acid-base indicator to find the end point — the point where the indicator changes colour, showing the acid and alkali have completely reacted. (p.170)
• For neutralisation of a mol of acid A with b mol of alkali B: M_aV_a / M_bV_b = a/b. (p.172)

6.8 Salts, Crystals and Their Uses in Daily Life

• Salt is an ionic compound formed when the H⁺ ion in an acid is replaced by a metal ion or ammonium ion, NH₄⁺. (p.175)
• Salt crystals have flat surfaces, straight sides, sharp vertices, a specific geometric shape and a fixed angle between adjacent surfaces. (p.176)

6.9 Preparation of Salts

• Solubility rule of thumb: all nitrate, sodium, potassium and ammonium salts are soluble; most sulphates are soluble except PbSO₄, BaSO₄, CaSO₄; most chlorides are soluble except Hg₂Cl₂, PbCl₂, AgCl; most carbonates are insoluble except Na₂CO₃, K₂CO₃, (NH₄)₂CO₃. (p.179)
• Soluble ammonium/sodium/potassium salts are prepared by acid + alkali (titration); other soluble salts by reacting excess reactive metal, metal oxide or metal carbonate with acid, then filtering, evaporating and crystallising. Insoluble salts are prepared by double decomposition (precipitation) of two soluble salt solutions. (p.180)
• In double decomposition, ions exchange between two aqueous salt solutions to form an insoluble precipitate and a new soluble salt, e.g. Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s). (p.187)

6.10 Effect of Heat on Salts

• Most carbonate salts decompose on heating to give a metal oxide + carbon dioxide gas (turns limewater cloudy); sodium and potassium carbonates do not decompose. (p.196)
• Nitrate salts decompose on heating to give a metal oxide + nitrogen dioxide gas (brown, acidic) + oxygen gas (relights a glowing splinter); sodium and potassium nitrates do not release the brown gas. (p.196)
• The colour of the residue on heating can indicate the cation present: yellow when hot/white when cool → Zn²⁺; brown when hot/yellow when cool → Pb²⁺; black (hot and cool) → Cu²⁺. (p.196)

6.11 Qualitative Analysis

• Qualitative analysis of a salt follows: observation of physical properties → effect of heat on the salt → anion and cation tests → confirmatory tests. (p.197)
• Anion tests: carbonate → add dilute acid, gas turns limewater cloudy; chloride → add dilute HNO₃ then AgNO₃(aq), white precipitate; sulphate → add dilute HCl then BaCl₂(aq), white precipitate; nitrate → add dilute H₂SO₄ and FeSO₄(aq), then concentrated H₂SO₄ down the side of the tube, a brown ring forms. (p.201)
• Cation tests use NaOH(aq) and NH₃(aq): the colour of any precipitate formed, and whether it dissolves in excess alkali, indicates which cation (e.g. Zn²⁺/Al³⁺/Pb²⁺ precipitates dissolve in excess NaOH; Cu²⁺ precipitate dissolves in excess NH₃ to a dark blue solution). (p.205)
• Confirmatory tests: NH₄⁺ with Nessler reagent (brown precipitate); Fe²⁺ with potassium hexacyanoferrate(III) (dark blue precipitate); Fe³⁺ with potassium thiocyanate (blood-red solution) or hexacyanoferrate(II) (dark blue precipitate); Al³⁺ vs Pb²⁺ distinguished using potassium iodide solution (Pb²⁺ gives a yellow precipitate soluble in hot water). (p.210)