Practise this lesson
Four printable worksheets that build from the foundations up to exam-style questions — start at whatever level suits you.
In 1856, photographer William Henry Fox Talbot patented the first photographic fixing process using sodium thiosulfate to dissolve excess silver chloride (AgCl) precipitate from photographic paper. His process depended on knowing which silver salts precipitate and which dissolve — the same solubility rules you will learn today. Three pairs of ionic solutions are mixed. Predict whether a precipitate forms in each case and, if so, name the precipitate:
- Potassium chloride + silver nitrate
- Potassium iodide + lead nitrate
- Sodium sulfate + barium nitrate
Write your predictions now — you may use any prior knowledge from Module 3. At the end of this lesson you will have the systematic rules to predict any precipitation reaction.
Know
- The NAGSAG solubility rules framework
- The three NESA-specified precipitation reactions
- How to write balanced molecular, complete ionic, and net ionic equations
Understand
- Why the four-ion method reliably predicts precipitation
- The difference between molecular and net ionic equations and when each is appropriate
- How precipitation reactions remove heavy metals in water treatment
Skills
- Predict precipitation from any combination of four ions using solubility rules
- Write balanced molecular and net ionic equations for precipitation reactions
- Apply precipitation chemistry to environmental and industrial contexts
Rather than memorising the solubility of every possible ionic compound individually, a small set of rules covers the vast majority of cases — ordered by priority to resolve any conflicts.
All soluble — no exceptions
All NH₄⁺ compounds soluble — no exceptions
Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺ — all soluble, no exceptions
Mostly soluble; except BaSO₄, PbSO₄ (insoluble); CaSO₄, Ag₂SO₄ (sparingly soluble)
CH₃COO⁻ — all soluble, no exceptions
CO₃²⁻, PO₄³⁻, OH⁻, S²⁻ — insoluble unless paired with Group 1 or NH₄⁺.
Cl⁻, Br⁻, I⁻: mostly soluble except Ag⁺ and Pb²⁺
NAGSAG solubility rules (priority order): N = Nitrates all soluble; A = Ammonium all soluble; G = Group 1 metals all soluble; S = Sulfates mostly soluble (exceptions: BaSO₄, PbSO₄ insoluble; CaSO₄, Ag₂SO₄ sparingly); A = Acetates all soluble; G = Generally insoluble (CO₃²⁻, PO₄³⁻, OH⁻, S²⁻ unless with Group 1 or NH₄⁺). Halide exceptions: AgCl/AgBr/AgI and PbCl₂/PbBr₂/PbI₂ all insoluble. Priority: N, A, G1 override all.
Write the NAGSAG framework with exceptions into your notes before the check below.
According to NAGSAG, all nitrate compounds are soluble in water with no exceptions.
We just saw the NAGSAG framework with its priority rules and key exceptions. That raises a question: when two solutions are mixed, how do you systematically combine NAGSAG with the actual ions present to predict whether a precipitate forms? This card answers it → by introducing the four-ion method and working through the KCl + AgNO₃ example to produce all three equation types.
Every precipitation prediction follows the same four-step procedure — identify all ions, determine possible products, apply solubility rules, write the equation.
Step 1: Identify the four ions present when two solutions are mixed.
Step 2: Identify the two possible new ionic combinations (each cation paired with the other anion).
Step 3: Apply NAGSAG to each possible product.
Step 4: Write the molecular, full ionic, and net ionic equations.
Applied to KCl(aq) + AgNO₃(aq):
- Ions present: K⁺, Cl⁻ (from KCl) and Ag⁺, NO₃⁻ (from AgNO₃)
- Possible products: AgCl (Ag⁺ + Cl⁻) and KNO₃ (K⁺ + NO₃⁻)
- NAGSAG: AgCl — Cl⁻ with Ag⁺ → INSOLUBLE. KNO₃ — Group 1 + NO₃⁻ → SOLUBLE
Molecular: $\text{KCl}(aq) + \text{AgNO}_3(aq) \rightarrow \text{AgCl}(s) + \text{KNO}_3(aq)$
Full ionic: $\text{K}^+(aq) + \text{Cl}^-(aq) + \text{Ag}^+(aq) + \text{NO}_3^-(aq) \rightarrow \text{AgCl}(s) + \text{K}^+(aq) + \text{NO}_3^-(aq)$
Net ionic: $\text{Ag}^+(aq) + \text{Cl}^-(aq) \rightarrow \text{AgCl}(s)$ (K⁺ and NO₃⁻ are spectator ions)
Four-ion precipitation method: (1) list the four ions from both solutions; (2) identify two possible new ionic pairings; (3) apply NAGSAG to each pairing — the insoluble one is the precipitate; (4) write molecular → full ionic (split all (aq) species, leave precipitate as formula(s)) → net ionic (cancel spectator ions). Never split the precipitate into ions.
Write the four-step method and the KCl + AgNO₃ three-equation example into your notes before the check below.
In the net ionic equation for a precipitation reaction, spectator ions are included on both sides.
Use what you have learned to solve the problems below. Show your reasoning clearly.
Question A: Explain how NAGSAG solubility rules can be applied in a real-world water treatment context.
Question B: Compare and contrast the four-ion method vs simply memorising known precipitates, using specific examples.
Question C: Predict what would happen if all nitrate compounds had exceptions to their solubility, and justify your prediction using evidence from the lesson.
We just saw the four-ion method for predicting precipitation using NAGSAG. That raises a question: which specific reactions does NESA require you to know completely — including colours, balanced equations, and net ionic forms? This card answers it → by presenting the three mandatory reactions (AgCl, PbI₂, BaSO₄) as a study table with all required details.
NESA specifies three precipitation reactions for Module 5 — know these completely, including precipitate formula, colour, balanced molecular equation, and net ionic equation.
| Reaction | Precipitate | Colour | Net ionic equation |
|---|---|---|---|
| KCl + AgNO₃ | AgCl(s) | White (curdy) | Ag⁺(aq) + Cl⁻(aq) → AgCl(s) |
| KI + Pb(NO₃)₂ | PbI₂(s) | Bright yellow ★ | Pb²⁺(aq) + 2I⁻(aq) → PbI₂(s) |
| Na₂SO₄ + Ba(NO₃)₂ | BaSO₄(s) | White | Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s) |
Balanced molecular equations:
$\text{KCl}(aq) + \text{AgNO}_3(aq) \rightarrow \text{AgCl}(s) + \text{KNO}_3(aq)$
$2\text{KI}(aq) + \text{Pb(NO}_3)_2(aq) \rightarrow \text{PbI}_2(s) + 2\text{KNO}_3(aq)$
$\text{Na}_2\text{SO}_4(aq) + \text{Ba(NO}_3)_2(aq) \rightarrow \text{BaSO}_4(s) + 2\text{NaNO}_3(aq)$
Three NESA precipitation reactions: (1) KCl + AgNO₃ → AgCl(s) white; net ionic: Ag⁺ + Cl⁻ → AgCl(s). (2) 2KI + Pb(NO₃)₂ → PbI₂(s) bright yellow; net ionic: Pb²⁺ + 2I⁻ → PbI₂(s). (3) Na₂SO₄ + Ba(NO₃)₂ → BaSO₄(s) white; net ionic: Ba²⁺ + SO₄²⁻ → BaSO₄(s). Always derive formula from charge balance — Pb²⁺ needs 2I⁻ → PbI₂, not PbI.
Copy the three-reaction table (reagents, precipitate, colour, net ionic equation) into your notes before the check below.
Which of the three NESA precipitation reactions produces a bright yellow precipitate?
We just saw the three NESA-mandated reactions with their balanced equations and net ionic forms. That raises a question: for any new precipitation reaction not in the memorised set, how do you construct all three equation types from scratch without making the common errors? This card answers it → by giving a five-step procedure and working through the Na₂SO₄ + Ba(NO₃)₂ example with a charge-balance check.
The three equations for a precipitation reaction are three levels of detail about the same event. Knowing all three shows complete understanding.
5-Step Procedure:
- Identify the precipitate using NAGSAG
- Write the balanced molecular equation: both (aq) reactants, precipitate (s), soluble product (aq)
- Write the full ionic equation: split all (aq) ionic compounds into ions; leave precipitate as formula (s)
- Identify and cancel spectator ions — ions appearing identically on both sides
- Write the net ionic equation; verify charge balance on both sides
Applied to Na₂SO₄ + Ba(NO₃)₂:
Molecular: $\text{Na}_2\text{SO}_4(aq) + \text{Ba(NO}_3)_2(aq) \rightarrow \text{BaSO}_4(s) + 2\text{NaNO}_3(aq)$
Full ionic: $2\text{Na}^+(aq) + \text{SO}_4^{2-}(aq) + \text{Ba}^{2+}(aq) + 2\text{NO}_3^-(aq) \rightarrow \text{BaSO}_4(s) + 2\text{Na}^+(aq) + 2\text{NO}_3^-(aq)$
Spectator ions: 2Na⁺ and 2NO₃⁻. Cancel them:
Net ionic: $\text{Ba}^{2+}(aq) + \text{SO}_4^{2-}(aq) \rightarrow \text{BaSO}_4(s)$
Charge check: left = +2 + (−2) = 0; right = 0 (neutral solid) ✓
Five-step net ionic equation procedure: (1) identify precipitate via NAGSAG; (2) write balanced molecular equation with state symbols; (3) full ionic = split all (aq) ionic compounds into ions, leave precipitate as formula(s); (4) cancel spectator ions (appear identically both sides); (5) write net ionic equation and verify charge balance. Key: never split the precipitate into ions.
Write the five-step procedure and the Na₂SO₄ + Ba(NO₃)₂ example with charge-balance check into your notes.
When writing a full ionic equation, an insoluble precipitate should be split into its constituent ions with an (aq) state symbol.
We just saw the five-step procedure for constructing molecular, full ionic, and net ionic equations from scratch. That raises a question: how is this precipitation chemistry applied in a real industrial context — and what decisions must an engineer make when choosing which reagent to add to contaminated water? This card answers it → by applying NAGSAG logic to the selection of precipitants for removing Pb²⁺ and Hg²⁺ from drinking water.
The same precipitation chemistry is applied daily in water treatment plants to remove dissolved lead, mercury, and arsenic from drinking water — turning invisible dissolved toxins into filterable solids.
Heavy metal contamination: lead from corroded pipes, mercury from industrial discharge, arsenic from geological sources. Standard treatment: add a reagent that forms an insoluble compound with the target ion, then filter or sediment the precipitate.
| Target ion | Precipitant added | Precipitate formed | Net ionic equation |
|---|---|---|---|
| Pb²⁺ | Na₂CO₃ | PbCO₃(s) — insoluble | Pb²⁺ + CO₃²⁻ → PbCO₃(s) |
| Pb²⁺ | Na₂SO₄ | PbSO₄(s) — insoluble | Pb²⁺ + SO₄²⁻ → PbSO₄(s) |
| Hg²⁺ | Na₂S | HgS(s) — black, extremely insoluble | Hg²⁺ + S²⁻ → HgS(s) |
The choice of precipitant must be specific — it must form an insoluble compound with the target metal but NOT with the other ions in the water (Na⁺, Ca²⁺, Mg²⁺, Cl⁻, etc.). This is an applied NAGSAG decision: which anion forms an insoluble compound with the target cation while remaining soluble with all other cations present?
Water treatment precipitation: Pb²⁺ removed by Na₂CO₃ (→ PbCO₃(s)) or Na₂SO₄ (→ PbSO₄(s)); Hg²⁺ by Na₂S (→ HgS(s) black, extremely insoluble). Precipitant selection rule: choose an anion that forms an insoluble compound with the target heavy-metal cation AND soluble compounds with all other cations present (Na⁺, Ca²⁺, Mg²⁺, etc.). Always verify both conditions using NAGSAG.
Write the precipitant selection logic and the three heavy-metal removal examples into your notes before the check below.
A precipitant for water treatment must form an insoluble compound with the target heavy metal ion, and should form soluble compounds with all other ions present in the water.
Predict whether a precipitate forms when each pair of solutions is mixed. If a precipitate forms, name it, state its colour, write the balanced molecular equation, and write the net ionic equation. (a) Na₂CO₃(aq) + CaCl₂(aq). (b) KNO₃(aq) + NaCl(aq). (c) FeCl₃(aq) + NaOH(aq).
Ions: Na⁺, CO₃²⁻ and Ca²⁺, Cl⁻. Possible products: CaCO₃ (Ca²⁺ + CO₃²⁻) and NaCl (Na⁺ + Cl⁻).
NAGSAG: CaCO₃ — CO₃²⁻ generally insoluble; Ca²⁺ is not Group 1 → INSOLUBLE. NaCl — Na⁺ Group 1 → soluble.
Precipitate: CaCO₃ (white).
Molecular: $\text{Na}_2\text{CO}_3(aq) + \text{CaCl}_2(aq) \rightarrow \text{CaCO}_3(s) + 2\text{NaCl}(aq)$
Net ionic: $\text{Ca}^{2+}(aq) + \text{CO}_3^{2-}(aq) \rightarrow \text{CaCO}_3(s)$ Charge check: +2 + (−2) = 0 ✓
Ions: K⁺, NO₃⁻ and Na⁺, Cl⁻. Possible products: KCl (K⁺ Group 1 → soluble) and NaNO₃ (Na⁺ Group 1, NO₃⁻ all soluble → soluble).
No precipitate forms. Both possible products are soluble — mixing these solutions produces no observable change.
Ions: Fe³⁺, Cl⁻ and Na⁺, OH⁻. Possible products: Fe(OH)₃ (Fe³⁺ + 3OH⁻) and NaCl (Na⁺ + Cl⁻).
NAGSAG: Fe(OH)₃ — OH⁻ generally insoluble; Fe³⁺ not Group 1 → INSOLUBLE. NaCl → soluble.
Precipitate: Fe(OH)₃ (rust-brown gelatinous precipitate).
Molecular: $\text{FeCl}_3(aq) + 3\text{NaOH}(aq) \rightarrow \text{Fe(OH)}_3(s) + 3\text{NaCl}(aq)$
Net ionic: $\text{Fe}^{3+}(aq) + 3\text{OH}^-(aq) \rightarrow \text{Fe(OH)}_3(s)$ Charge check: +3 + 3(−1) = 0 ✓
A water sample contains Pb²⁺ (5.0 × 10⁻³ mol/L) along with Na⁺, Ca²⁺, Cl⁻, and NO₃⁻. (a) Identify two anions that would precipitate Pb²⁺. (b) Evaluate which is preferable. (c) Write the net ionic equation for the preferred reaction.
From NAGSAG — Pb²⁺ forms insoluble compounds with: SO₄²⁻ (PbSO₄ — insoluble) and CO₃²⁻ (PbCO₃ — generally insoluble). Adding Na₂SO₄ or Na₂CO₃ would precipitate Pb²⁺.
SO₄²⁻ (Na₂SO₄): Na⁺ + SO₄²⁻ → Na₂SO₄ — soluble ✓. But Ca²⁺ + SO₄²⁻ → CaSO₄ — sparingly soluble. Adding excess Na₂SO₄ might cause partial CaSO₄ precipitation. Problem: Ca²⁺ may be unintentionally removed.
CO₃²⁻ (Na₂CO₃): Na⁺ + CO₃²⁻ → Na₂CO₃ — soluble ✓. But Ca²⁺ + CO₃²⁻ → CaCO₃ — insoluble. Na₂CO₃ would co-precipitate Ca²⁺ as well.
Preferred: Na₂CO₃ efficiently precipitates Pb²⁺ (more decisively insoluble than CaSO₄). If Ca²⁺ removal is acceptable, Na₂CO₃ is the better choice. If Ca²⁺ must remain, Na₂SO₄ is preferable as CaSO₄ is only sparingly soluble and its precipitation is less complete.
$\text{Pb}^{2+}(aq) + \text{CO}_3^{2-}(aq) \rightarrow \text{PbCO}_3(s)$
Charge check: +2 + (−2) = 0 = 0 ✓
- Define the term "precipitate" and explain its importance in predicting ionic reactions.
- Identify one common misconception about solubility rules and precipitation reactions. Explain why it is incorrect.
- Describe how the NAGSAG rules relate to the nitrate ion using an example from the lesson.
Q1. Which of the following pairs of solutions will produce a precipitate when mixed?
Q2. When KI(aq) is mixed with Pb(NO₃)₂(aq), a bright yellow precipitate forms. Which is the correct net ionic equation?
Q3. A student writes the full ionic equation for Na₂CO₃(aq) + CaCl₂(aq) as: 2Na⁺(aq) + CO₃²⁻(aq) + Ca²⁺(aq) + 2Cl⁻(aq) → Ca²⁺(aq) + CO₃²⁻(aq) + 2Na⁺(aq) + 2Cl⁻(aq). What error has the student made?
SAQ 1 (3 marks): Using the four-ion method, predict whether a precipitate forms when aqueous iron(III) chloride is added to aqueous sodium hydroxide. If a precipitate forms, name it, state its colour, and write the net ionic equation including a charge balance check.
SAQ 2 (4 marks): A water treatment engineer needs to remove Pb²⁺ ions from contaminated groundwater that also contains Na⁺, Ca²⁺, Mg²⁺, Cl⁻, and NO₃⁻. Evaluate whether Na₂SO₄ or Na₂CO₃ would be the better precipitant to use, applying NAGSAG solubility rules in your reasoning. Write the net ionic equation for the preferred reaction.
Show Answers
SAQ 1: Ions present: Fe³⁺, Cl⁻ (from FeCl₃) and Na⁺, OH⁻ (from NaOH). Possible products: Fe(OH)₃ (Fe³⁺ + 3OH⁻) and NaCl (Na⁺ + Cl⁻). NAGSAG: Fe(OH)₃ — OH⁻ generally insoluble, Fe³⁺ not Group 1 → INSOLUBLE. NaCl — Na⁺ Group 1 → soluble. Precipitate: Fe(OH)₃, rust-brown gelatinous solid. Net ionic: Fe³⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(s). Charge check: +3 + 3(−1) = 0 = 0 ✓
SAQ 2: SO₄²⁻: PbSO₄ is insoluble ✓ but CaSO₄ is sparingly soluble — may co-precipitate Ca²⁺. CO₃²⁻: PbCO₃ is insoluble ✓ but CaCO₃ is also insoluble — would remove Ca²⁺ as well. Na₂CO₃ precipitates Pb²⁺ more completely (PbCO₃ is decisively insoluble). If Ca²⁺ removal is acceptable, Na₂CO₃ is preferred. Net ionic: Pb²⁺(aq) + CO₃²⁻(aq) → PbCO₃(s). Charge check: +2 + (−2) = 0 ✓
Put the method for predicting whether a precipitate forms when two solutions are mixed in the correct order.
- Calculate Qsp using the initial (diluted) ion concentrations.
- Identify all ions present after mixing the two solutions.
- State whether a precipitate forms and write the net ionic equation.
- Write the possible ionic compounds and look up their solubility / Ksp.
- Compare Qsp to Ksp: if Qsp > Ksp, a precipitate forms.
Complete the Learn phase to unlock practice questions.
A student is investigating the precipitation of heavy metal ions from contaminated industrial wastewater. The sample contains Pb²⁺ (8.0 × 10⁻³ mol/L), Fe³⁺ (5.0 × 10⁻³ mol/L), Na⁺, Cl⁻, and NO₃⁻. The student plans to add NaOH solution followed by Na₂SO₄ solution to selectively remove the heavy metal ions. Evaluate this plan using NAGSAG solubility rules: predict which precipitates will form at each step, write net ionic equations for each precipitation, and assess whether the sequence is appropriate or could be improved.
Look back at your Think First predictions about KCl + AgNO₃, KI + Pb(NO₃)₂, and Na₂SO₄ + Ba(NO₃)₂. Recall Fox Talbot's 1856 photographic fixing process — he needed to know exactly which silver salts precipitate. Were you correct in your predictions? What do you now understand about using NAGSAG systematically to predict precipitation, the way Talbot needed to predict which salts would dissolve or precipitate in his developing process?