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Module 3 · L3 of 12 35 min ⚡ +50 XP in Learn · +25 to complete

Precipitation & Solubility Rules

In 2014, the city of Flint, Michigan switched its water supply and, because the new water was not treated for corrosion control, it began corroding old lead pipes and dissolving Pb²⁺ ions into people’s tap water. The long-term fix was not to filter lead out of the water but to stop it dissolving in the first place: dosing the water with orthophosphate so it builds a protective mineral scale on the inside of the pipes, alongside replacing the lead pipes. The solubility chemistry behind that protective coating is what this lesson is about.

Today's hook, In 2014, the city of Flint, Michigan switched its water supply and, because the new water was not treated for corrosion control, it began corroding old lead pipes and dissolving Pb²⁺ ions into people’s tap water. The long-term fix was not to filter lead out of the water but to stop it dissolving in the first place: dosing the water with orthophosphate so it builds a protective mineral scale on the inside of the pipes, alongside replacing the lead pipes. The solubility chemistry behind that protective coating is what this lesson is about.
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Warm up and recall

Warm up first

Three quick questions from earlier lessons. Pulling old material back to mind before you learn something new makes the new material stick better, so this is not busywork.

Worksheets

Practise this lesson

Four printable worksheets that build from the foundations up to exam-style questions, start at whatever level suits you.

Build Foundations & guided practice Apply Application practice Master Mastery challenge Exam Exam-style questions
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Recall, your gut answer first
+5 XP warm-up

A water treatment plant receives water contaminated with dissolved lead ions from old pipes. The lead is invisible, the water looks perfectly clear. Engineers add a chemical to the water and within seconds a white solid forms and sinks to the bottom, taking the lead with it.

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What you'll master, and the words for it

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What you'll master
Know

Key facts

  • The NAGSAG solubility rules and their exceptions
  • Common precipitates and their colours
  • What spectator ions are
Understand

Concepts

  • Why mixing two clear solutions can produce a solid
  • The difference between molecular, full ionic, and net ionic equations
  • How precipitation is used in water treatment
Can do

Skills

  • Apply solubility rules to predict if a precipitate forms
  • Write molecular, full ionic, and net ionic equations
  • Identify spectator ions in a reaction
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Key terms
Precipitation reaction
Two aqueous solutions react to form an insoluble solid (precipitate): A(aq) + B(aq) → AB(s) + C(aq).
Solubility rules (NAGSAG)
Mnemonic for soluble groups: Nitrates, Ammonium salts, Group 1 salts, Sulfates (except Ba²⁺, Pb²⁺, Ca²⁺), Acetates, Group 17 salts (except Ag⁺, Pb²⁺, Hg₂²⁺).
Net ionic equation
An equation showing only the ions that participate in the reaction, omitting spectator ions.
Spectator ions
Ions that are present in solution but do not participate in the net reaction; appear on both sides and cancel out.
Precipitate
An insoluble solid that forms when two ionic solutions are mixed and the product ions combine to form a compound that is insoluble according to the solubility rules.
Double displacement
Two ionic compounds exchange anions in solution; occurs if one product is insoluble, a gas, or a weak electrolyte.
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Solubility Rules, The NAGSAG Framework

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Solubility Rules, The NAGSAG Framework
core concept

When two clear solutions are mixed and a white or coloured solid instantly appears, you have just witnessed a precipitation reaction, two ions met in solution and found they could not stay dissolved together. You can predict which combinations do this before mixing, using one tool: the NAGSAG solubility rules.

NAGSAG solubility rules as a comparison table with 6 ion-group rows.
Ion Solubility Insoluble Exceptions
NO₃⁻ (nitrate)All solubleNone
NH₄⁺ (ammonium)All solubleNone
Group 1 metal ionsAll solubleNone
SO₄²⁻ (sulfate)Mostly solubleBa²⁺, Pb²⁺, Ca²⁺
Cl⁻, Br⁻, I⁻ (halides)Mostly solubleAg⁺, Pb²⁺
CO₃²⁻ (carbonate)Mostly insolubleGroup 1, NH₄⁺
OH⁻ (hydroxide)Mostly insolubleGroup 1, Ba²⁺
Apply rules to both possible products: When two solutions are mixed, check both possible product combinations separately, one may be soluble, the other insoluble.
Calcium sulfate exception: CaSO₄ is insoluble, students frequently forget this because calcium compounds are often soluble. Memorise it: Ca²⁺ is an exception for sulfates.
Medical insight, BaSO₄: Barium sulfate is so insoluble it is used as a "barium meal" in medical X-rays of the digestive system. Patients swallow it without absorbing any into the bloodstream, its insolubility is what makes it safe as a contrast agent.

NAGSAG solubility rules (a simplified course prediction table, use the NESA data sheet in exams): Nitrates, Ammonium salts, Group 1 salts, Acetates, all soluble (no common exceptions); Sulfates, mostly soluble except Ba²⁺, Pb²⁺, Ca²⁺; halides mostly soluble except Ag⁺, Pb²⁺; carbonates, hydroxides, sulfides mostly insoluble except with Group 1 or NH₄⁺. Some salts have further exceptions (e.g. certain fluorides), so treat the table as a guide, not an absolute law.

Pause, copy the highlighted rule into your book before moving on.

Explain it: Using the NAGSAG rules, predict whether a precipitate forms when aqueous lead(II) nitrate Pb(NO₃)₂ is mixed with aqueous potassium iodide KI. Name the precipitate and explain your reasoning. (2–3 sentences)

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Writing Full Ionic and Net Ionic Equations

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Writing Full Ionic and Net Ionic Equations
core concept

We just saw that NAGSAG rules let us predict which ionic combinations form precipitates. That raises a question: how do we write equations that show exactly which ions react and which are bystanders? This card answers it → the three levels of ionic equation, from molecular to net ionic.

Safety, lead compounds, teacher-controlled. Lead(II) salts such as lead nitrate and lead iodide are toxic and a cumulative poison. The lead iodide colour is best shown as a teacher demonstration, or a strictly supervised small-scale activity: wear safety glasses and gloves, use minimum volumes of dilute solution, never pipette by mouth, wash hands afterwards, and collect ALL lead-containing waste in a labelled heavy-metal waste container, never down the drain. The same solubility rule can be shown more safely with silver chloride or barium sulfate.

There are three levels of equation for precipitation reactions. Using the reaction of Pb(NO₃)₂(aq) with KI(aq) as the example:

Level 1, Molecular equation (complete formulas):
Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq)
Level 2, Full ionic equation (split all soluble ionic compounds into ions; keep precipitate intact):
Pb²⁺(aq) + 2NO₃⁻(aq) + 2K⁺(aq) + 2I⁻(aq) → PbI₂(s) + 2K⁺(aq) + 2NO₃⁻(aq)
Spectator ions appear identically on both sides, K⁺(aq) and NO₃⁻(aq) here. They play no role in the reaction.
Level 3, Net ionic equation (remove spectator ions, the most chemically meaningful):
Pb²⁺(aq) + 2I⁻(aq) → PbI₂(s)
Rule for ionic equations: Only split compounds into ions if they are aqueous (aq) AND ionic. Never split a precipitate (s), a liquid (l), or a gas (g) into ions.
Critical error, splitting the precipitate: Students frequently write Pb²⁺(aq) + 2I⁻(aq) on the right side. The precipitate is insoluble, it exists as a solid lattice, not free ions. Always keep the precipitate as PbI₂(s).

The net ionic equation removes spectator ions, showing only the ions that react: e.g. Pb²⁺(aq) + 2I⁻(aq) → PbI₂(s). Never split the precipitate into ions, it is an insoluble solid lattice, not free aqueous ions.

Add the highlighted rule to your notes before the check below.

Match it: Match each equation type to its correct description for the reaction between Pb(NO₃)₂(aq) and KI(aq).

  • Molecular equation
  • Full ionic equation
  • Net ionic equation
  • Spectator ions
  • K⁺(aq) and NO₃⁻(aq), appear unchanged on both sides
  • Pb²⁺(aq) + 2I⁻(aq) → PbI₂(s)
  • Pb²⁺(aq) + 2NO₃⁻(aq) + 2K⁺(aq) + 2I⁻(aq) → PbI₂(s) + 2K⁺(aq) + 2NO₃⁻(aq)
  • Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq)
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Water Treatment, Precipitation Removing Heavy Metals

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Water Treatment, Precipitation Removing Heavy Metals
core concept

We just saw how to write net ionic equations that isolate exactly which ions react to form a precipitate. That raises a question: where are these precipitation reactions actually used in the real world? This card answers it → water engineers use the same reactions to remove toxic heavy metal ions from contaminated water supplies.

The precipitation reactions you've just learned are the same reactions municipal water engineers use to remove dissolved toxic metals from contaminated water supplies.

Heavy metal ions such as Pb²⁺, Hg²⁺, and Cd²⁺ dissolve from corroded pipes and industrial runoff. They are colourless and tasteless, contaminated water looks identical to clean water. Treatment adds a reagent that triggers precipitation:

Pb²⁺

Reagent added: Na₂SO₄(aq)
Precipitate formed: PbSO₄(s), white
Net ionic equation: Pb²⁺(aq) + SO₄²⁻(aq) → PbSO₄(s)

Pb²⁺

Reagent added: Na₂CO₃(aq)
Precipitate formed: PbCO₃(s), white
Net ionic equation: Pb²⁺(aq) + CO₃²⁻(aq) → PbCO₃(s)

Hg²⁺

Reagent added: Na₂S(aq)
Precipitate formed: HgS(s), black
Net ionic equation: Hg²⁺(aq) + S²⁻(aq) → HgS(s)
💧 Real-World Anchor, Water Treatment: In short answer Q3, you will be asked to select an appropriate reagent to remove a specific ion from water. Use solubility rules to choose a reagent whose anion forms an insoluble compound with the target cation, and a soluble compound with all other ions present.
Practical limitation: Adding excess reagent can cause secondary contamination, for example, too much Na₂S introduces sulfide ions into the water. Real-world applications require precise stoichiometry. In HSC questions, note this limitation if asked to evaluate the method.

To remove a heavy metal ion from water, add a reagent whose anion forms an insoluble compound with the target cation, e.g. add Na₂S to precipitate Hg²⁺ as HgS(s). The precipitate is removed by filtration; excess reagent must be controlled to avoid secondary contamination.

Pause, write the highlighted point into your book.

Fill the gap: To remove dissolved Hg²⁺ ions from contaminated water, engineers add Na₂S solution. The Hg²⁺ ions combine with S²⁻ ions to form [___](s), a black insoluble solid that can be removed by filtration.

Cross-lesson links: In L01–L02 you classified chemical changes and saw synthesis/decomposition patterns, precipitation is a third type where dissolved ions exchange partners. In L05, you will apply similar prediction logic to acid-carbonate reactions, where CO₂ gas (rather than a solid precipitate) signals the reaction.
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Worked examples

Worked examples · reveal as you go

Worked Example 1, Predicting Whether a Precipitate Forms +5 XP on full reveal

Predict whether a precipitate will form when aqueous barium chloride (BaCl₂) is mixed with aqueous sodium sulfate (Na₂SO₄). If a precipitate forms, name it, give its colour, and write the molecular equation with state symbols.

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BaCl₂ provides: Ba²⁺ and Cl⁻   |   Na₂SO₄ provides: Na⁺ and SO₄²⁻
Identify all ions present when solutions are mixed
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BaSO₄ (Ba²⁺ + SO₄²⁻)   and   NaCl (Na⁺ + Cl⁻)
Identify the two possible product combinations
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NaCl: Group 1 (Na⁺), all soluble ✓   |   BaSO₄: sulfate, with Ba²⁺, insoluble exception → precipitate forms
Apply solubility rules to each product
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BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)
Atom check: Left, 1 Ba, 2 Cl, 2 Na, 1 S, 4 O. Right, 1 Ba, 2 Cl, 2 Na, 1 S, 4 O. ✓
Write and balance the molecular equation
Step 1 of 4
Worked Example 2, Writing the Net Ionic Equation +5 XP on full reveal

For the reaction in Worked Example 1, write (a) the full ionic equation and (b) the net ionic equation. Identify the spectator ions.

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Ba²⁺(aq) + 2Cl⁻(aq) + 2Na⁺(aq) + SO₄²⁻(aq) → BaSO₄(s) + 2Na⁺(aq) + 2Cl⁻(aq)
Write the full ionic equation, split soluble compounds; keep BaSO₄(s) intact
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Na⁺(aq) and Cl⁻(aq) appear identically on both sides, they are spectator ions.
Identify spectator ions
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Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)
Cancel spectator ions to give the net ionic equation
Step 1 of 3
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What you'll master, and the words for it

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Formula reference · this lesson
core formula
📐

Key Tools, This Lesson

$\text{X}^{n+}_{(aq)} + \text{Y}^{m-}_{(aq)} \rightarrow \text{XY}_{(s)}\downarrow$  (net ionic)
Solubility rules (NAGSAG): used to predict whether an ionic compound will dissolve (soluble) or precipitate (insoluble) Net ionic equation: full ionic equation minus spectator ions, ions that appear identically on both sides Precipitation indicator: cloudy or coloured solid ($\downarrow$) appearing in a previously clear solution
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Common misconception

Common errors · the 3 traps that cost marks

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Common misconception

All ionic compounds dissolve in water because water is polar.

Fix: Water dissolves many ionic compounds through ion-dipole interactions, but not all. Solubility depends on the balance between lattice energy and hydration energy. Compounds with very high lattice energy (e.g., AgCl, BaSO₄) are insoluble despite water's polarity.

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Splitting the precipitate into ions in the ionic equation

When writing the full ionic equation, students split all compounds into ions including the precipitate, producing something like Ba²⁺(aq) + SO₄²⁻(aq) → Ba²⁺(aq) + SO₄²⁻(aq).

Fix: Only dissolve (aq) species into their ions. The precipitate, labelled (s), exists as an insoluble ionic lattice, not free ions in solution, so it is never split. Write the precipitate formula intact: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s). The net ionic equation is then just this line, there are no spectator ions in this example.

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Cancelling ions that are not truly identical on both sides

Students cancel an ion from the full ionic equation even when it appears in a different state or formula on each side (e.g. one side has Cu²⁺(aq) and the other has Cu(s)).

Fix: Spectator ions must appear with the same formula, charge, and state (aq) on both sides before they can be cancelled. If an ion changes state (e.g. Cu²⁺(aq) is reduced to Cu(s)), it is NOT a spectator, it is a participant. Cancel only ions that are completely unchanged on both sides of the full ionic equation.

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Drill, then revisit

Quick-fire practice · 5 reps +2 XP per reveal

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Aqueous silver nitrate AgNO₃(aq) is mixed with aqueous sodium chloride NaCl(aq). Does a precipitate form?

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Aqueous potassium carbonate K₂CO₃(aq) is mixed with aqueous calcium chloride CaCl₂(aq). Does a precipitate form?

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Aqueous sodium sulfate Na₂SO₄(aq) is mixed with aqueous potassium nitrate KNO₃(aq). Does a precipitate form?

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Q1 (4 marks): Explain what a spectator ion is and why it does not appear in the net ionic equation. Use the reaction between potassium iodide and lead(II) nitrate as your example. Write the full ionic equation and identify the spectator ions.

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Q2 (4 marks): Predict and describe what happens when aqueous iron(III) chloride (FeCl₃) is mixed with aqueous sodium hydroxide (NaOH). (a) Use solubility rules to predict whether a precipitate forms, and identify it with its colour. (1 mark) (b) Write the balanced molecular equation with state symbols. (1 mark) (c) Write the net ionic equation. (1 mark) (d) Identify the spectator ions. (1 mark)

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Revisit your thinking

Earlier you read about Flint, Michigan (2014), where Pb²⁺ ions dissolved from corroding pipes contaminated the tap water. Now you can explain the chemistry of the fix: utilities add orthophosphate, which reacts with lead at the pipe surface to build a very insoluble lead-phosphate scale (for example 3Pb²⁺ + 2PO₄³⁻ → Pb₃(PO₄)₂(s)). This protective coating forms on the pipe wall and lowers how much lead dissolves into the water, rather than precipitating and filtering lead out of every glass; replacing the lead pipes was also essential. Here insolubility is used to keep a metal locked in a solid scale, not to filter a bulk precipitate.

This works because when two ionic solutions mix, every possible ion combination is instantly tested against the solubility rules, whichever combinations are insoluble precipitate immediately. You can predict this before mixing, which is exactly what water engineers do when choosing a reagent.

Now revisit your initial response. What did you get right? What has changed in your thinking?

Look back at your initial response in your book. Annotate it with what you now understand differently.

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Interactive Tool, Reaction Types Explorer Open fullscreen ↗
The Reaction Types Explorer tool shows a precipitation reaction. Which of the following correctly describes what happens to spectator ions during the reaction?
✓

Multiple choice

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Multiple choice
+5 XP per correct · +25 XP all-correct

Pick your answer, then rate your confidence. That tells the system what to drill next.

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Short answer
UnderstandBand 34 MARKS

Q1. Explain what a spectator ion is and why it does not appear in the net ionic equation. Use the reaction between potassium iodide and lead(II) nitrate as your example. Write the full ionic equation and identify the spectator ions.

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ApplyBand 44 MARKS

Q2. Predict and describe what happens when aqueous iron(III) chloride (FeCl₃) is mixed with aqueous sodium hydroxide (NaOH). (a) Use solubility rules to predict whether a precipitate forms, and identify it with its colour. (1 mark) (b) Write the balanced molecular equation with state symbols. (1 mark) (c) Write the net ionic equation. (1 mark) (d) Identify the spectator ions. (1 mark)

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EvaluateBand 55 MARKS

Q3. A water treatment facility receives water contaminated with dissolved mercury(II) ions (Hg²⁺). An engineer proposes adding sodium sulfide (Na₂S) to the water to remove the mercury. (a) Write the net ionic equation for the removal of Hg²⁺ using Na₂S. (1 mark) (b) Use solubility rules to justify why mercury(II) sulfide (HgS) precipitates. (1 mark) (c) Evaluate one advantage and one limitation of using precipitation to remove heavy metals from water supplies. (2 marks) (d) The engineer adds excess Na₂S to ensure complete mercury removal. Assess the risk this creates. (1 mark)

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📖 Comprehensive answers (click to reveal)

Activity 1, Predicting Precipitates

1. AgNO₃ + NaCl: Ions: Ag⁺, NO₃⁻, Na⁺, Cl⁻. Products: AgCl (Ag⁺ + Cl⁻, insoluble, white precipitate), NaNO₃ (soluble). Net ionic: Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

2. K₂CO₃ + CaCl₂: Ions: K⁺, CO₃²⁻, Ca²⁺, Cl⁻. Products: CaCO₃ (insoluble, white precipitate), KCl (Group 1 → soluble). Net ionic: Ca²⁺(aq) + CO₃²⁻(aq) → CaCO₃(s)

3. Na₂SO₄ + KNO₃: Ions: Na⁺, SO₄²⁻, K⁺, NO₃⁻. All possible products (NaNO₃, K₂SO₄) involve Group 1 → all soluble. No precipitate forms.

Activity 2, Water Treatment Table

Row 1 (Pb²⁺, with Na⁺ Cl⁻): Add Na₂CO₃ or Na₂SO₄. With Na₂CO₃: PbCO₃(s) white precipitate forms. Net ionic: Pb²⁺(aq) + CO₃²⁻(aq) → PbCO₃(s)

Row 2 (Cu²⁺, with Na⁺ NO₃⁻): Add NaOH. Cu(OH)₂(s) pale blue precipitate forms (Cu²⁺ + OH⁻; copper is not Group 1 so Cu(OH)₂ insoluble). Net ionic: Cu²⁺(aq) + 2OH⁻(aq) → Cu(OH)₂(s)

Row 3 (Fe³⁺, with K⁺ SO₄²⁻): Add KOH or NaOH. Fe(OH)₃(s) rust brown precipitate (Fe³⁺ + OH⁻; iron not Group 1 so Fe(OH)₃ insoluble). Net ionic: Fe³⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(s)

Why not NaCl for Pb²⁺? The Cl⁻ ion is already present in solution. Adding NaCl increases the Cl⁻ concentration but PbCl₂ is only slightly soluble (not completely insoluble like PbSO₄ or PbCO₃), significant Pb²⁺ would remain in solution. Additionally, the Cl⁻ ion already present means no driving force for additional precipitation from NaCl.

❓ Multiple Choice

PbI₂: Pb²⁺ is an exception for iodides → insoluble. The others: Na₂CO₃ (Group 1 → soluble), (NH₄)₂SO₄ (ammonium → soluble), KOH (Group 1 → soluble).

The precipitate is AgCl(s). Na⁺ and NO₃⁻ remain dissolved unchanged, spectator ions.

CaCO₃ is an insoluble ionic solid (not free ions). It must be written as CaCO₃(s) in any ionic equation.

Cu²⁺ + 2OH⁻ → Cu(OH)₂(s). Spectator ions Na⁺ and SO₄²⁻ are removed. Options A and B are molecular/full ionic equations, not net ionic. Option C is incorrect (CuSO₄ is soluble).

A white solid forming and persisting (cannot be redissolved) confirms new substance PbSO₄(s) produced, strongest evidence of chemical change. Temperature change alone is insufficient; cloudiness could be physical.

6. C (Band 5) PbCl₂ is only slightly soluble (it sits at the borderline of the solubility exceptions), meaning Pb²⁺ concentration would not drop to safe levels. Also, Cl⁻ ions added to drinking water have taste/health implications.

7. B (Band 6) Both BaSO₄ and CaSO₄ are insoluble sulfates (both are exceptions to the solubility rules), so Na₂SO₄ would precipitate both. Na₂CO₃ would also precipitate both (BaCO₃ and CaCO₃ are insoluble). BaSO₄ is substantially less soluble than CaSO₄, so at carefully controlled reagent concentrations, BaSO₄ precipitates preferentially, but this precision is not expected at Year 11 level. Option B is the best available answer at this level.

Short Answer Model Answers

Q8 (4 marks): A spectator ion is an ion that appears unchanged (same formula, same state) on both sides of the full ionic equation, it does not participate in the reaction [1]. It is excluded from the net ionic equation because it plays no role in forming the product [1]. Full ionic equation for Pb(NO₃)₂ + 2KI: Pb²⁺(aq) + 2NO₃⁻(aq) + 2K⁺(aq) + 2I⁻(aq) → PbI₂(s) + 2K⁺(aq) + 2NO₃⁻(aq) [1]. Spectator ions: K⁺(aq) and NO₃⁻(aq), they appear unchanged on both sides [1].

Q9 (4 marks): (a) Fe(OH)₃, rust brown precipitate. Fe³⁺ with OH⁻: iron is not Group 1 → hydroxide insoluble; Na⁺ with Cl⁻ → NaCl soluble [1]. (b) FeCl₃(aq) + 3NaOH(aq) → Fe(OH)₃(s) + 3NaCl(aq) [1, balanced with state symbols]. (c) Fe³⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(s) [1]. (d) Spectator ions: Na⁺(aq) and Cl⁻(aq) [1].

Q10 (5 marks): (a) Hg²⁺(aq) + S²⁻(aq) → HgS(s) [1]. (b) HgS is a sulfide with a non-Group 1, non-ammonium cation → falls under "generally insoluble" rule for sulfides; very low solubility is consistent with precipitation [1]. (c) Advantage: converts dissolved invisible ions into filterable solids, removing them effectively and at relatively low cost [1]. Limitation: not all heavy metal compounds are sufficiently insoluble to reduce contamination below safe limits, or may require precise stoichiometry [1]. (d) Excess Na₂S introduces S²⁻ ions into the treated water, which is itself toxic/harmful and constitutes secondary contamination, negating the purpose of treatment [1].