Acids, bases, salts and titrations

The pH scale and indicators

The pH scale runs from 0 to 14. A pH of 7 is neutral; below 7 is acidic; above 7 is alkaline. Each unit change in pH represents a 10-fold change in hydrogen ion concentration.

• pH 0–6 → acidic (excess H⁺ ions)
• pH 7 → neutral
• pH 8–14 → alkaline (excess OH⁻ ions)

Universal indicator changes colour across the pH range (red in strong acids → purple in strong alkalis). A pH meter gives a more precise numerical reading.

Common laboratory indicators:

• Litmus: red in acid, blue in alkali
• Phenolphthalein: colourless in acid, pink/red in alkali — used in titrations where alkali is in the burette
• Methyl orange: red in acid, yellow in alkali — used where acid is added from burette

Acids and alkalis

Acids ionise in water to produce H⁺ (or H₃O⁺) ions:

• Hydrochloric acid: HCl(aq) → H⁺(aq) + Cl⁻(aq)
• Sulfuric acid: H₂SO₄(aq) → 2H⁺(aq) + SO₄²⁻(aq)
• Nitric acid: HNO₃(aq) → H⁺(aq) + NO₃⁻(aq)

Strong acids fully ionise; weak acids partially ionise (e.g. ethanoic acid, carbonic acid).

Alkalis are bases that dissolve in water. They produce OH⁻ ions:

• Sodium hydroxide: NaOH(aq) → Na⁺(aq) + OH⁻(aq)
• Potassium hydroxide: KOH

Neutralisation

Acid + base → salt + water

General equations:

• Acid + metal oxide → salt + water
• Acid + metal hydroxide → salt + water
• Acid + metal carbonate → salt + water + carbon dioxide
• Acid + metal → salt + hydrogen

The ionic equation for all neutralisation reactions: H⁺(aq) + OH⁻(aq) → H₂O(l)

Naming salts

The name of a salt has two parts: the cation (from the base, usually a metal) and the anion (from the acid).

Salt preparation methods

Method 1 — Titration (soluble salt from soluble acid + soluble base): e.g. NaOH + HCl → NaCl + H₂O Use indicator to find exact volumes, then repeat without indicator, evaporate to crystallise.

Method 2 — Insoluble base + acid (soluble salt without carbonate fizz): e.g. excess CuO powder + dilute H₂SO₄, warm, filter off excess CuO, evaporate filtrate to crystallise CuSO₄·5H₂O.

Method 3 — Metal + acid (reactive metals): e.g. excess Zn + H₂SO₄ → ZnSO₄ + H₂, filter excess Zn, evaporate.

Method 4 — Precipitation (insoluble salt): Mix two soluble solutions whose ions form an insoluble product: Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq) Filter the precipitate, wash, dry.

Acid–base titrations

A titration precisely measures the volume of one solution needed to neutralise a known volume of another.

Procedure (WJEC required practical):

1. Pipette a fixed volume of alkali into a conical flask using a volumetric pipette.
2. Add a few drops of indicator (e.g. phenolphthalein).
3. Fill burette with acid. Record initial burette reading.
4. Add acid slowly until the indicator just changes colour permanently (the end point).
5. Record final burette reading. Calculate titre = final − initial.
6. Repeat until concordant results (within 0.10 cm³).
7. Use the mean of concordant titres to calculate concentration.

Calculation: n(acid) = c(acid) × V(acid) From balanced equation, find n(base) c(base) = n(base) ÷ V(base)

Common examiner traps

1. Using the wrong indicator: phenolphthalein changes colour at pH ~8–10 (good for strong base + weak acid titrations); methyl orange at pH ~3–4 (good for strong acid + weak base).
2. Forgetting to use mean of concordant titres: identify and discard anomalous results before averaging.
3. Not washing the conical flask between runs (with distilled water — does not affect moles of alkali).
4. Confusing salt names: nitric acid → nitrate (NOT nitrite); sulfuric acid → sulfate (NOT sulfite).
5. Carbonate reaction: acid + carbonate gives CO₂ gas (effervescence) — don't forget the third product.