Solubility Rules

Summary: All sodium, potassium, and ammonium salts are soluble. All nitrates are soluble. Most chlorides are soluble (except AgCl, PbCl₂). Most sulfates are soluble (except BaSO₄, PbSO₄, CaSO₄). Most carbonates and hydroxides are insoluble. Tags: igcse chemistry acids-bases Created: 2026-07-14 Last Updated: 2026-07-16

Definition of Solubility

Solubility is the mass of solute that dissolves in 100 g of water at a given temperature to form a saturated solution. It is usually expressed in grams per 100 g of water.

A saturated solution is one in which no more solute can dissolve at that temperature — any additional solute added will remain undissolved. A precipitate is an insoluble solid that forms when two solutions are mixed and an insoluble compound is produced.

The solubility of most solids increases with temperature, though the exact relationship varies from substance to substance.

Solubility rules allow chemists to predict whether a salt will dissolve in water. These rules are essential for:

• Predicting whether a precipitate will form when two solutions are mixed
• Choosing the correct method of salt preparation
• Writing net ionic equations for precipitation reactions

The Solubility Rules

Always Soluble

All sodium (Na⁺) salts are soluble.

• NaCl, NaNO₃, Na₂SO₄, Na₂CO₃ — all dissolve readily in water

All potassium (K⁺) salts are soluble.

• KCl, KNO₃, K₂SO₄, K₂CO₃ — all dissolve readily in water

All ammonium (NH₄⁺) salts are soluble.

• NH₄Cl, NH₄NO₃, (NH₄)₂SO₄, (NH₄)₂CO₃ — all dissolve readily in water

All nitrates (NO₃⁻) are soluble.

• No exceptions. Every metal nitrate dissolves in water.

These four rules have zero exceptions and are the most reliable starting point for any solubility question.

Mostly Soluble

Chlorides (Cl⁻)

Most chlorides are soluble. The key exceptions are:

• Silver chloride (AgCl) — white precipitate, insoluble in water and dilute acids
• Lead(II) chloride (PbCl₂) — sparingly soluble in cold water, soluble in hot water. Crystallises out on cooling.

Bromides and iodides follow the same pattern as chlorides.

Sulfates (SO₄²⁻)

Most sulfates are soluble. The key exceptions are:

• Barium sulfate (BaSO₄) — white precipitate, insoluble
• Lead(II) sulfate (PbSO₄) — white precipitate, insoluble
• Calcium sulfate (CaSO₄) — sparingly soluble

Sparingly soluble means a small amount dissolves — enough to create a slightly saturated solution, but much less than a fully soluble salt.

Mostly Insoluble

Carbonates (CO₃²⁻)

Most carbonates are insoluble. The only exceptions are:

• Sodium carbonate (Na₂CO₃) — soluble
• Potassium carbonate (K₂CO₃) — soluble
• Ammonium carbonate ((NH₄)₂CO₃) — soluble

Every other carbonate — including CaCO₃ (limestone), MgCO₃, CuCO₃, ZnCO₃, FeCO₃ — is insoluble.

Hydroxides (OH⁻)

Most hydroxides are insoluble. The exceptions are:

• Sodium hydroxide (NaOH) — soluble
• Potassium hydroxide (KOH) — soluble
• Ammonium hydroxide (NH₄OH) — soluble (formed when ammonia gas dissolves in water)
• Calcium hydroxide (Ca(OH)₂) — sparingly soluble (limewater)

Oxides (O²⁻)

Most oxides are insoluble except for Group 1 metal oxides (Na₂O, K₂O) which react with water to form soluble hydroxides.

Summary Table of Solubility Rules

Predicting Precipitate Formation

Solubility rules are used to predict whether a precipitate will form when two aqueous solutions are mixed (a double displacement / precipitation reaction).

The method:

1. Write the formulae of the two reactants
2. Swap the positive and negative ions to determine the two possible products
3. Check the solubility of each product using the solubility rules
4. If either product is insoluble, it will appear as a precipitate
5. Write the state symbols: (aq) for soluble products in solution, (s) for the precipitate

Example: “What happens when aqueous silver nitrate is mixed with aqueous sodium chloride?”

• Reactants: AgNO₃(aq) + NaCl(aq)
• Possible products (swap ions): AgCl and NaNO₃
• Check solubility: AgCl is insoluble (exception to chloride rule), NaNO₃ is soluble (all nitrates soluble)
• Result: AgCl forms as a white precipitate
• Equation: AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)

Example: “What happens when aqueous barium chloride is mixed with aqueous sodium sulfate?”

• Reactants: BaCl₂(aq) + Na₂SO₄(aq)
• Possible products (swap ions): BaSO₄ and NaCl
• Check solubility: BaSO₄ is insoluble (exception to sulfate rule), NaCl is soluble
• Result: BaSO₄ forms as a white precipitate
• Equation: BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)

Example: “Will a precipitate form when aqueous potassium nitrate is mixed with aqueous sodium chloride?”

• Reactants: KNO₃(aq) + NaCl(aq)
• Possible products: KCl and NaNO₃
• Check solubility: KCl is soluble (all potassium salts soluble), NaNO₃ is soluble (all nitrates soluble)
• Result: No precipitate — both products are soluble, all ions remain in solution

Using Solubility Rules for Salt Preparation

Solubility rules determine which method of salt preparation is appropriate:

Preparing an insoluble salt — Precipitation method:

• Mix two soluble salts, each containing one of the required ions
• The insoluble salt forms as a precipitate
• Filter, wash with distilled water, and dry

Example: To prepare barium sulfate (BaSO₄):

• Choose two soluble salts: BaCl₂(aq) — soluble chloride, Na₂SO₄(aq) — soluble sulfate
• Mix: BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)
• Filter off the white BaSO₄ precipitate

Preparing a soluble salt:

• If the salt is soluble, precipitation will not work (ions stay in solution)
• Use one of: acid + metal, acid + base, acid + carbonate, or titration (for Group 1 / ammonium salts)
• See Making Salts for full details

Ionic Equations for Precipitation Reactions

In a precipitation reaction, the driving force is the combination of two ions to form an insoluble solid. Spectator ions (those that remain aqueous on both sides) can be cancelled to write a net ionic equation.

Example: AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)

Full ionic equation (all aqueous ionic compounds split into ions):

Ag⁺(aq) + NO₃⁻(aq) + Na⁺(aq) + Cl⁻(aq) → AgCl(s) + Na⁺(aq) + NO₃⁻(aq)

Cancel spectator ions (Na⁺ and NO₃⁻ appear unchanged on both sides):

Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

Net ionic equation: Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

Example: BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)

Full ionic equation:

Ba²⁺(aq) + 2Cl⁻(aq) + 2Na⁺(aq) + SO₄²⁻(aq) → BaSO₄(s) + 2Na⁺(aq) + 2Cl⁻(aq)

Cancel spectator ions (Na⁺ and Cl⁻):

Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)

Net ionic equation: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)

Test for Sulfate Ions (SO₄²⁻)

This is a key practical application of solubility rules:

Method:

1. Add a few drops of barium chloride solution (BaCl₂) to the test solution
2. Add a few drops of dilute hydrochloric acid (HCl)
3. A white precipitate of barium sulfate confirms the presence of sulfate ions

Why add HCl?

• If a white precipitate forms without HCl, it could be BaCO₃ or BaSO₃ (both white and insoluble) rather than BaSO₄
• Adding acid removes carbonate and sulfite ions by converting them to gases:
  • CO₃²⁻ + 2H⁺ → CO₂(g) + H₂O
  • SO₃²⁻ + 2H⁺ → SO₂(g) + H₂O
• Barium sulfate is insoluble in acid, so the precipitate that remains after acid addition confirms sulfate

Ionic equation: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)

Test for Halide Ions (Cl⁻, Br⁻, I⁻)

This uses the insolubility of silver halides:

Method:

1. Add a few drops of silver nitrate solution (AgNO₃) to the test solution
2. Add a few drops of dilute nitric acid (HNO₃)
3. Observe the colour of the precipitate

Ionic equation (for chloride): Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

Key Concepts from Past Papers

• substances required:
• steam / water
• conditions: one mark each for any two of:
• 6000 kPa / 60 atm (pressure)
• acid catalyst
• salt carbon dioxide water
• add (hydrochloric) acid
• test gas given off with limewater

Keywords from Past Papers

acid, add, aqueous, precipitate, indicator, carbon, dioxide, white, sodium, hydroxide, colour, nitrogen, limewater, temperature, point

Related Notes

• Acids and Bases
• Indicators and pH
• Making Salts
• Neutralization

Sources

• OpenStax Chemistry 2e, Chapter 11 (Solutions and Colloids), Chapter 15 (Equilibria of Other Reaction Classes) — openstax.org/books/chemistry-2e/
• BBC Bitesize GCSE Chemistry, Solubility and Precipitates — bbc.co.uk/bitesize/topics/
• Cambridge IGCSE Chemistry 0620 Syllabus, Topic 12: Acids, Bases and Salts
• CK-12 Chemistry, Solubility Rules and Precipitation Reactions — ck12.org/chemistry/

Past Paper Sources

• 0620/31 May/June 2015: Q11(c)(i) (0m), Q66(a)(i) (0m), Q66(a)(ii) (0m)
• 0620/32 Feb/March 2015: Q66(e)(ii) (2m), Q66(b)(ii) (1m), Q77(d)(i) (0m)
• 0620/32 Feb/March 2017: Q22(f)(ii) (1m), Q44(d)(i) (1m), Q55(f)(i) (1m) (+1 more)
• 0620/32 Feb/March 2018: Q88(b)(ii) (3m), Q22(a)(iv) (2m), Q77(b)(i) (1m) (+1 more)
• 0620/32 Feb/March 2019: Q44(b)(ii) (1m)
• 0620/32 Feb/March 2020: Q55(d)(iii) (2m), Q66(a)(ii) (1m)
• 0620/32 Feb/March 2021: Q66(a)(ii) (1m)
• 0620/32 Feb/March 2022: Q88(d)(ii) (2m), Q44(a)(ii) (0m)
• 0620/32 Feb/March 2023: Q33(b)(ii) (1m), Q88(c)(iii) (1m)
• 0620/32 May/June 2018: Q33(b)(i) (1m), Q33(b)(ii) (1m)
• 0620/32 May/June 2020: Q88(e)(i) (1m)
• 0620/33 May/June 2016: Q88(c)(ii) (1m), Q88(c)(iii) (0m)

Common Misconceptions