Acids, Bases and Salt Preparations

Salt preparation choice depends on TWO questions:

Question 1: Is the TARGET SALT soluble or insoluble?

Question 2: Are the REACTANTS soluble or insoluble?

Decision tree.

                       Is target salt soluble?
                                 |
              +------------------+------------------+
              | YES                                  | NO
              v                                      v
   Is there an INSOLUBLE                  PRECIPITATION
   reactant available?                    Mix two soluble
              |                            solutions whose
   +----------+----------+                 ions form
   | YES                  | NO            insoluble salt
   v                      v               → filter, wash, dry
ACID + INSOLUBLE       TITRATION
BASE/METAL/CARBONATE    (acid + alkali)
+ filter + crystallise

Method 1: Acid + insoluble base / metal / carbonate (soluble salt + insoluble reactant).

Used for salts like CuSO₄, ZnCl₂, MgSO₄, FeCl₂. The base/metal/carbonate is INSOLUBLE in water (so excess can be filtered off).

• Add reactant in EXCESS to warm acid; stir.
• Stop when no more reactant dissolves.
• Filter to remove unreacted excess.
• Evaporate the filtrate to saturated solution.
• Cool slowly for crystals; filter; wash with cold distilled water; pat dry.

Why excess? Guarantees ALL the acid has been neutralised — no acid contamination of the salt.

Method 2: Titration (soluble salt + two soluble reactants).

Used for sodium / potassium / ammonium salts (no insoluble reactant available because Group 1 and NH₄⁺ salts are all soluble).

• Titrate acid against alkali to find the exact volume needed (using indicator).
• Note the volume of acid required (the 'titre').
• Repeat WITHOUT indicator: mix the same volume of acid and alkali in a clean flask (no indicator → no contamination of crystals).
• Evaporate to saturated solution.
• Cool slowly to crystallise; filter; wash; dry.

Method 3: Precipitation (insoluble salt).

Used for AgCl, BaSO₄, PbI₂.

• Pick TWO SOLUBLE solutions whose ions combine to give the insoluble salt:
  • Want AgCl? → Use AgNO₃(aq) + NaCl(aq) (both soluble).
  • Want BaSO₄? → Use BaCl₂(aq) + Na₂SO₄(aq).
  • Want PbI₂? → Use Pb(NO₃)₂(aq) + KI(aq).
• Mix the solutions → insoluble salt precipitates immediately.
• Filter to collect precipitate as residue.
• Wash with distilled water (removes soluble byproducts like NaNO₃).
• Dry in a warm oven or air-dry.

Worked procedure — CuSO₄·5H₂O from CuO + dilute H₂SO₄.

Equation: $\text{CuO(s)} + \text{H}_2\text{SO}_4\text{(aq)} \rightarrow \text{CuSO}_4\text{(aq)} + \text{H}_2\text{O(l)}$.

Why this method? Cu is BELOW H — Cu metal doesn't react with dilute H₂SO₄ (no fizz). CuO is an INSOLUBLE BASE — can be filtered off. CuSO₄ is SOLUBLE — can crystallise from solution.

Step-by-step.

1. Measure ~50 cm³ of dilute H₂SO₄ into a 100 cm³ beaker. Warm gently with a Bunsen burner + tripod + gauze (or on a hot-water bath) to ~50 °C. Warming speeds the reaction. Do NOT boil.

2. Add black CuO powder a spatula at a time, stirring after each. The CuO will gradually dissolve, giving a blue solution of CuSO₄. Continue adding until NO MORE CuO DISSOLVES — this is the visual sign that ALL the acid has reacted. Adding CuO IN EXCESS guarantees no leftover acid (otherwise it would crystallise with the salt).

3. Filter to remove excess CuO. Set up a filter funnel with filter paper in a clean conical flask (or evaporating basin). Pour the warm mixture through. Unreacted CuO is RESIDUE on the filter paper; CuSO₄ solution is the FILTRATE.

4. Evaporate the filtrate to a saturated solution. Heat gently in an evaporating basin on a water bath (or low Bunsen flame). Watch for the solution becoming thicker / cloudier. Test for saturation: dip a clean glass rod into the solution; if blue crystals form on the rod as it dries, the solution is saturated. STOP heating here — DO NOT evaporate to dryness (would lose water of crystallisation, decompose to anhydrous CuSO₄).

5. Cool slowly to crystallise. Leave to cool to room temperature undisturbed (ideally overnight). Slow cooling = large, well-formed CuSO₄·5H₂O crystals (blue). The solubility decreases as the solution cools, so crystals come out of solution.

6. Filter, wash, dry. Filter the crystals from the remaining 'mother liquor'. Wash with a SMALL volume of COLD distilled water (cold so crystals don't redissolve; small to minimise yield loss; distilled to avoid contamination). Pat dry between two sheets of filter paper.

Final product: pure, dry, blue CuSO₄·5H₂O crystals (hydrated copper(II) sulfate pentahydrate).

Variations.

• Use ZnO + H₂SO₄ → ZnSO₄·7H₂O (white crystals).
• Use MgO + 2HCl → MgCl₂·6H₂O (colourless).
• Use FeO + H₂SO₄ → FeSO₄·7H₂O (pale green).
• Replace insoluble base with insoluble carbonate (e.g. ZnCO₃ + H₂SO₄ → ZnSO₄ + H₂O + CO₂) — same method.
• Replace insoluble base with reactive metal (e.g. Zn + H₂SO₄ → ZnSO₄ + H₂) — note H₂ gas.

Insoluble salts (AgCl, BaSO₄, PbI₂, PbCl₂) cannot be made by 'acid + base + crystallise' methods. Instead, they are made by PRECIPITATION — mixing two soluble solutions whose ions combine to form the insoluble salt.

Worked example — PbI₂ ('golden rain').

Equation: $\text{Pb(NO}_3)_2\text{(aq)} + 2\text{KI(aq)} \rightarrow \text{PbI}_2\text{(s)} + 2\text{KNO}_3\text{(aq)}$.

Choose the reactants:

• Need Pb²⁺: use Pb(NO₃)₂ (all nitrates soluble). NOT PbCl₂ or PbSO₄ — those are insoluble.
• Need I⁻: use KI (all K salts soluble). NaI also OK; AgI is NOT (AgI insoluble).
• The other ions (K⁺, NO₃⁻) must form a SOLUBLE compound (KNO₃) so they wash away in the filtrate.

Step-by-step.

1. Mix equal volumes of Pb(NO₃)₂(aq) and KI(aq) in a beaker. A bright yellow precipitate of PbI₂ forms immediately — described as 'golden rain' because of the shimmering appearance.

2. Filter through filter paper. The yellow PbI₂ stays as RESIDUE on the filter; the colourless KNO₃ solution passes through as FILTRATE (discarded).

3. Wash the residue ON the filter paper with several portions of distilled water. This removes any soluble impurities (K⁺, NO₃⁻, excess Pb²⁺ or I⁻ from one reactant being in slight excess). Doesn't dissolve PbI₂ — it's insoluble.

4. Dry. Transfer to a watch glass and either dry in a warm oven or leave at room temperature.

Final product: pure, dry, yellow PbI₂ powder.

Other precipitation examples.

Salt	Soluble reactants	Other product (soluble)
AgCl (white)	AgNO₃ + NaCl	NaNO₃
AgBr (cream)	AgNO₃ + NaBr	NaNO₃
AgI (yellow)	AgNO₃ + NaI	NaNO₃
BaSO₄ (white)	BaCl₂ + Na₂SO₄	NaCl
PbSO₄ (white)	Pb(NO₃)₂ + Na₂SO₄	NaNO₃
PbCl₂ (white)	Pb(NO₃)₂ + NaCl	NaNO₃

Ionic equation (omits spectator ions).

For PbI₂: $\text{Pb}^{2+}\text{(aq)} + 2\text{I}^-\text{(aq)} \rightarrow \text{PbI}_2\text{(s)}$.

The K⁺ and NO₃⁻ are spectator ions — they appear on both sides of the full equation and don't actually take part in the precipitate-forming step.

Edexcel-required solubility rules — memorise these for both calculation and method-choice questions.

Always SOLUBLE.

Group of salts	Notes
All Group 1 (Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺)	No exceptions
All ammonium (NH₄⁺)	No exceptions
All nitrates (NO₃⁻)	No exceptions — including AgNO₃, Pb(NO₃)₂, Cu(NO₃)₂, Fe(NO₃)₃
All ethanoates (CH₃COO⁻)	(Not always required by 4CH1 but useful)

MOSTLY soluble, with exceptions:

Salt type	Soluble	INSOLUBLE exceptions
Chlorides (Cl⁻)	Most metal chlorides	AgCl (white), PbCl₂ (white, soluble in HOT water)
Bromides (Br⁻)	Most metal bromides	AgBr (cream), PbBr₂ (white)
Iodides (I⁻)	Most metal iodides	AgI (yellow), PbI₂ (yellow)
Sulfates (SO₄²⁻)	Most metal sulfates	BaSO₄ (white), PbSO₄ (white), CaSO₄ (slightly)

MOSTLY insoluble, with exceptions:

Salt type	Insoluble	SOLUBLE exceptions
Carbonates (CO₃²⁻)	Most metal carbonates	Group 1 (Li₂CO₃, Na₂CO₃, K₂CO₃) + (NH₄)₂CO₃
Hydroxides (OH⁻)	Most metal hydroxides	Group 1 (LiOH, NaOH, KOH) + Ca(OH)₂ slightly + NH₃ solution
Sulfides (S²⁻)	Most metal sulfides	Group 1 + Group 2 + (NH₄)₂S

Practical use in method choice.

• Want CuSO₄ → SOLUBLE (most sulfates). Use insoluble-base method.
• Want NaCl → SOLUBLE (Group 1). Use titration method.
• Want AgCl → INSOLUBLE (exception). Use precipitation.
• Want BaSO₄ → INSOLUBLE (exception). Use precipitation.
• Want PbI₂ → INSOLUBLE (exception). Use precipitation.

Memory tricks.

• 'SAN-N' — Sodium, Ammonium, Nitrate (always soluble) + 'and Nitrate again' to emphasise.
• 'AgCl, AgBr, AgI' — all silver halides insoluble (the basis for the halide test).
• 'BaSO₄' — barium sulfate insoluble (the basis for the sulfate test).
• Carbonates 'all insoluble except K, Na, NH₄' (heat them → metal oxide + CO₂; but Group 1 carbonates are too stable to decompose easily).

Water of crystallisation = water molecules chemically bound into the crystal lattice of a salt. They form a regular part of the crystal structure — NOT 'free' water trapped in cracks, but specific H₂O units coordinated to the metal cation.

Formula notation. Write the salt formula, then a dot, then the number of H₂O molecules: $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$ means each formula unit of CuSO₄ has 5 water molecules associated with it.

Common hydrated salts.

Salt	Formula	Colour	Common name
Copper(II) sulfate pentahydrate	CuSO₄·5H₂O	Bright blue	'Blue vitriol'
Magnesium sulfate heptahydrate	MgSO₄·7H₂O	Colourless	Epsom salt
Sodium carbonate decahydrate	Na₂CO₃·10H₂O	White	Washing soda
Calcium sulfate dihydrate	CaSO₄·2H₂O	White	Gypsum
Iron(II) sulfate heptahydrate	FeSO₄·7H₂O	Pale green	Green vitriol
Cobalt(II) chloride hexahydrate	CoCl₂·6H₂O	Pink	(Cobalt chloride paper)

Heating drives off water — hydrated → anhydrous.

When a hydrated salt is heated, the water of crystallisation is driven off as steam, leaving the anhydrous (water-free) salt:

$\text{CuSO}_4 \cdot 5\text{H}_2\text{O(s)} \xrightarrow{\Delta} \text{CuSO}_4\text{(s)} + 5\text{H}_2\text{O(g)}$

Colour change observed: blue → white (for CuSO₄). Condensation forms on cold parts of the apparatus (the steam).

Reverse — adding water reforms the hydrate.

If water is added to anhydrous (white) CuSO₄, the blue colour returns and heat is given out (exothermic):

$\text{CuSO}_4\text{(s)} + 5\text{H}_2\text{O(l)} \rightarrow \text{CuSO}_4 \cdot 5\text{H}_2\text{O(s)} + \text{heat}$

Tests using hydrated/anhydrous salts.

• Anhydrous CuSO₄ test for water: white CuSO₄ turns BLUE on adding water. Used to detect water in unknown liquids.
• Cobalt(II) chloride paper test for water: blue CoCl₂ paper turns PINK on adding water.

Both confirm the PRESENCE of water but NOT purity (a salt solution would also turn the paper pink).

Why slow cooling matters for crystallisation. Hydrated salts crystallise out of solution with their full water of crystallisation when given time. Fast cooling gives small, irregular crystals; slow cooling gives large, well-formed crystals with the correct hydration state.