The Characteristic Properties of Acids and Bases

Acid. A substance that produces hydrogen ions ($\text{H}^{+}$) in aqueous solution.

• Common: $\text{HCl}$ (hydrochloric), $\text{H}_{2}\text{SO}_{4}$ (sulfuric), $\text{HNO}_{3}$ (nitric), $\text{CH}_{3}\text{COOH}$ (ethanoic / acetic).

Base. Any substance that neutralises an acid. Examples: metal oxides ($\text{CuO}$), metal hydroxides ($\text{NaOH}$), carbonates ($\text{CaCO}_{3}$), ammonia ($\text{NH}_{3}$).

Alkali. A SOLUBLE base. Produces hydroxide ions ($\text{OH}^{-}$) in aqueous solution.

• $\text{NaOH}, \text{KOH}, \text{Ca(OH)}_{2}, \text{NH}_{3}\text{(aq)}$.

Universal indicator and pH. A mixture of dyes that changes colour with H⁺ concentration.

pH	Colour	Acidic / Neutral / Alkaline
0-2	Red	Strong acid
3-4	Orange	Weak acid
5-6	Yellow	Slightly acidic
7	Green	Neutral
8-9	Light blue	Slightly alkaline
10-12	Blue	Weak alkali
13-14	Purple/violet	Strong alkali

Worked. Lemon juice (pH ≈ 3) turns universal indicator orange — slightly acidic.

1. Acid + reactive metal → salt + hydrogen. $\text{Mg} + 2\text{HCl} \to \text{MgCl}_{2} + \text{H}_{2}.$

• Test for hydrogen: lit splint, "squeaky pop".
• Reactive metals: anything ABOVE hydrogen in the reactivity series (Mg, Al, Zn, Fe).
• Cu and below DON'T react with dilute acids.

2. Acid + base/alkali → salt + water (neutralisation). $\text{HCl} + \text{NaOH} \to \text{NaCl} + \text{H}_{2}\text{O}.$

• Net ionic: $\text{H}^{+} + \text{OH}^{-} \to \text{H}_{2}\text{O}$.
• Same equation regardless of which acid + which alkali — that's the essence of neutralisation.

3. Acid + carbonate → salt + water + carbon dioxide. $\text{CaCO}_{3} + 2\text{HCl} \to \text{CaCl}_{2} + \text{H}_{2}\text{O} + \text{CO}_{2}.$

• Test for CO₂: limewater turns CLOUDY (white precipitate of CaCO₃).
• Reaction effervesces (fizzes from CO₂ gas).

4. Acid + ammonia → ammonium salt. $\text{HCl} + \text{NH}_{3} \to \text{NH}_{4}\text{Cl}.$

• Often used industrially: ammonium nitrate fertiliser, ammonium sulfate fertiliser.

Salt naming. The salt is named from the acid:

• HCl → chloride.
• HNO₃ → nitrate.
• H₂SO₄ → sulfate.
• $\text{CH}_{3}\text{COOH}$ → ethanoate.

Worked. Sulfuric acid + magnesium oxide → magnesium sulfate + water. $\text{H}_{2}\text{SO}_{4} + \text{MgO} \to \text{MgSO}_{4} + \text{H}_{2}\text{O}.$

Strong acid. Ionises COMPLETELY in water — every molecule releases H⁺.

• $\text{HCl}(aq) \to \text{H}^{+}(aq) + \text{Cl}^{-}(aq)$ (irreversible arrow).
• Examples: HCl, HNO₃, H₂SO₄.

Weak acid. PARTIALLY ionises — only a small fraction of molecules release H⁺. The rest remain as molecules; the system is at equilibrium.

• $\text{CH}_{3}\text{COOH}(aq) \rightleftharpoons \text{CH}_{3}\text{COO}^{-}(aq) + \text{H}^{+}(aq)$.
• Examples: ethanoic acid, citric acid, carbonic acid.

Same concentration, different acidity. $1\,\text{mol/dm}^{3}$ HCl has pH $\approx 0$ (high H⁺); $1\,\text{mol/dm}^{3}$ ethanoic acid has pH $\approx 2.4$ (much less H⁺ released).

Concentration ≠ strength.

• A DILUTE strong acid has pH $\sim 4$.
• A CONCENTRATED weak acid has pH $\sim 2$.
• The two ideas combine to determine pH.

Strong vs weak alkalis.

• Strong (fully ionises): NaOH, KOH.
• Weak (partial): NH₃(aq) — equilibrium with $\text{NH}_{4}^{+} + \text{OH}^{-}$.

Tests to compare. Equal concentrations:

• Strong acid: more vigorous reaction with magnesium (more H⁺), pH closer to 0.
• Same conductivity comparison: strong acids conduct better (more ions per volume).

Cambridge tip. Cambridge expects the explanation in terms of FULL vs PARTIAL ionisation, NOT in terms of "stronger / more concentrated" — those are different ideas.