State that bases are oxides or hydroxides of metals and that alkalis are soluble bases

Acids, Bases and Salts – Characteristic Properties (Cambridge IGCSE 0620)

Learning Objectives

By the end of this lesson you should be able to:

• Define acids, bases and alkalis exactly as the syllabus states.
• State that bases are metal oxides or metal hydroxides and that alkalis are the soluble bases.
• List the characteristic properties of acids and bases (five points each).
• Classify oxides as acidic, basic or amphoteric, linking the classification to the metallic or non‑metallic nature of the element.
• Write and balance typical acid‑base reactions, including neutralisation and gas‑evolving reactions, using correct state symbols.
• Explain the pH/pOH scale, convert between pH, [H⁺] and [OH⁻], and use a universal‑indicator colour table.
• Apply the full set of solubility rules (including all exceptions) to predict the formation of salts and precipitation.
• Describe the laboratory preparation, separation and purification of salts.

1. Definitions (Syllabus 7.1)

2. Characteristic Properties

2.1 Acids (five points)

1. Turn blue litmus paper red.
2. Taste sour (never taste in the laboratory!).
3. React with metals to give a salt and hydrogen gas.
   Example: Zn(s) + 2 HCl(aq) → ZnCl₂(aq) + H₂(g)
4. React with carbonates or bicarbonates to give a salt, water and carbon dioxide gas.
   Example: CaCO₃(s) + 2 HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)
5. Conduct electricity when aqueous (they are electrolytes).

2.2 Bases (including alkalis) – five points

1. Turn red litmus paper blue.
2. Feel slippery or soapy (saponification of skin oils).
3. React with acids to give a salt and water (neutralisation).
   Example: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)
4. Conduct electricity when aqueous.
5. Often have a bitter taste (never taste in the laboratory!).

3. Oxides – Acidic, Basic or Amphoteric? (Syllabus 7.2)

• Basic oxides are metal oxides. They react with water to give a base (metal hydroxide).
  Example: CaO(s) + H₂O(l) → Ca(OH)₂(aq)
• Acidic oxides are non‑metal oxides. They react with water to give an acid.
  Example: SO₃(g) + H₂O(l) → H₂SO₄(aq)
• Amphoteric oxides (e.g. Al₂O₃, ZnO) can behave either as an acid or a base depending on the reaction partner.

4. Solubility Rules (Syllabus 7.3)

5. Typical Acid‑Base Reactions

5.1 Neutralisation (acid + base → salt + water)

General form (state symbols included):

Acid(aq) + Base(aq) → Salt(aq) + Water(l)

Examples:

• HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)
• H₂SO₄(aq) + 2 NaOH(aq) → Na₂SO₄(aq) + 2 H₂O(l)
• HNO₃(aq) + Ca(OH)₂(aq) ⇌ Ca(NO₃)₂(aq) + 2 H₂O(l) (slightly soluble base, equilibrium shown)

5.2 Gas‑Evolving Reactions

• Acid + Metal → Salt + H₂(g)
  Mg(s) + 2 HCl(aq) → MgCl₂(aq) + H₂(g)
• Acid + Carbonate / Bicarbonate → Salt + H₂O + CO₂(g)
  NaHCO₃(s) + HCl(aq) → NaCl(aq) + H₂O(l) + CO₂(g)
• Acid + Sulphide → Salt + H₂S(g)
  FeS(s) + 2 HCl(aq) → FeCl₂(aq) + H₂S(g)

5.3 Precipitation of Insoluble Hydroxides

When a strong alkali reacts with a soluble metal salt, an insoluble metal hydroxide may precipitate:

Metal‑salt(aq) + Base(aq) → Metal‑hydroxide(s) + Salt‑of‑base(aq)

Example (copper(II) hydroxide precipitate):

6. pH, pOH and Universal Indicator (Syllabus 7.1)

• pH = –log₁₀[H⁺]  (0 ≤ pH ≤ 14 at 25 °C)
• pOH = –log₁₀[OH⁻]  and pH + pOH = 14
• Acidic solution: pH < 7 (high [H⁺], low [OH⁻])
• Neutral solution: pH = 7 (pure water)
• Basic solution: pH > 7 (low [H⁺], high [OH⁻]; alkalis give pH > 7)

Worked example – Convert pH = 3 to ion concentrations:

pH = –log[H⁺] ⇒ [H⁺] = 10⁻³ M
pOH = 14 – pH = 11 ⇒ [OH⁻] = 10⁻¹¹ M

Universal‑indicator colour chart (simplified)

Use the colour change to estimate the pH of an unknown solution – a common practical skill (AO 3).

7. Preparation of Salts (Syllabus 7.3)

1. Neutralisation method – acid + alkali → salt + water.
   Example: H₂SO₄(aq) + 2 NaOH(aq) → Na₂SO₄(aq) + 2 H₂O(l)
   Laboratory steps: Mix the solutions, evaporate the filtrate to dryness, cool to crystallise, filter and dry the solid salt.
2. Precipitation method – mix two soluble salts; an insoluble product precipitates (chosen using solubility rules).
   Example: BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2 NaCl(aq)
   Laboratory steps: Stir, allow the precipitate to settle, filter, wash with cold water, and dry.
3. Acid + metal – yields a soluble salt and H₂ gas; evaporate the solution to isolate the salt.
   Example: 2 HCl(aq) + Zn(s) → ZnCl₂(aq) + H₂(g)
4. Acid + carbonate – gives a soluble salt, water and CO₂; evaporate to obtain the solid salt.
   Example: 2 HNO₃(aq) + CaCO₃(s) → Ca(NO₃)₂(aq) + H₂O(l) + CO₂(g)

In each method the choice of reactants is guided by the solubility rules so that the desired salt remains in solution while any by‑product (e.g., a precipitate or gas) can be removed.

8. Summary Table – Acids vs. Bases (incl. Alkalis)

9. Suggested Diagram (Flow‑chart)

• Metal oxides → (add H₂O) → Metal hydroxides → (if soluble) → Alkalis
• Non‑metal oxides → (add H₂O) → Acids
• Acid + Alkali → Neutralisation → Salt + Water
• Acid + Metal / Carbonate → Salt + Gas (H₂ or CO₂)

10. Quick‑fire Revision Checklist

• Write the exact syllabus definitions of acid, base and alkali.
• State the five characteristic properties of acids and of bases.
• Identify whether a given oxide is acidic, basic or amphoteric, linking it to the metal / non‑metal character.
• Balance a neutralisation equation with correct state symbols.
• Convert a given pH to [H⁺] (and to pOH/[OH⁻] if required).
• Use the full set of solubility rules, including the sulphate and chloride exceptions, to predict precipitation.
• Describe the laboratory steps (reaction, filtration, evaporation/crystallisation) for preparing a salt by any of the four methods.