Difficulty: Introductory to Intermediate | Prerequisites: Ion charges, polyatomic ions, writing ionic compound formulas, balancing equations
Precipitation reactions are the practical payoff of learning to write balanced equations and knowing your ions. When two aqueous ionic solutions are mixed, the ions are free to recombine. Whether anything interesting happens depends entirely on the solubility rules: if one of the possible new pairings is insoluble, it crashes out of solution as a solid precipitate. This is fundamental to qualitative analysis in the lab, water treatment, and any situation where you need to remove a specific ion from solution.
Mix two salt solutions, swap the ions, and check solubility rules. If a new combination is insoluble, you get a precipitate. If everything stays dissolved, no reaction occurs.
Soluble
A substance that dissolves in water to form a homogeneous solution, dissociating into its component ions. In simple terms, it disappears into the water.
Insoluble
A substance that does not dissolve appreciably in water and remains as a solid. In simple terms, it refuses to dissolve and you can see it sitting there.
Precipitate
The insoluble solid product formed when two aqueous solutions are mixed and an insoluble compound results from the ion recombination. Think of it as the solid that "falls out" of the solution.
Molecular equation
The balanced equation showing the complete formulas of all reactants and products, including their states of matter.
Complete ionic equation
The equation in which all soluble strong electrolytes are written as dissociated ions, while insoluble compounds, gases, and water remain as complete formulas.
Net ionic equation
The equation that shows only the ions and molecules that actually participate in the reaction, with spectator ions removed. Think of it as the equation stripped down to just the bit that matters.
Spectator ions
Ions present in solution that do not participate in the reaction. They appear on both sides of the complete ionic equation and cancel out. In simple terms, they are in the audience, not on stage.
Tags: soluble, insoluble, precipitate, molecular equation, complete ionic equation, net ionic equation, spectator ions, aqueous, dissociation
You need these memorised. They tell you which ionic compounds dissolve in water and which do not.
Alkali metal salts (Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺): almost always soluble.
Ammonium salts (NH₄⁺): almost always soluble.
Nitrates (NO₃⁻): always soluble. No exceptions you need to know.
Acetates (CH₃COO⁻): almost always soluble.
Halides (Cl⁻, Br⁻, I⁻): generally soluble, except with Ag⁺, Pb²⁺, and Hg₂²⁺.
Sulfates (SO₄²⁻): generally soluble, except with Ba²⁺, Pb²⁺, Ca²⁺ (slightly), and Sr²⁺.
Hydroxides (OH⁻): insoluble, except with alkali metals, Ba²⁺, Sr²⁺, and Ca²⁺ (slightly).
Sulfides (S²⁻): insoluble, except with alkali metals, alkaline earth metals, and NH₄⁺.
Carbonates (CO₃²⁻): insoluble, except with alkali metals and NH₄⁺.
Phosphates (PO₄³⁻): insoluble, except with alkali metals and NH₄⁺.
The recitation asked which compounds are insoluble. The correct answers:
AgCl : halide with Ag⁺, insoluble.
AgI : halide with Ag⁺, insoluble.
PbCl₂ : halide with Pb²⁺, insoluble.
BaSO₄ : sulfate with Ba²⁺, insoluble.
Ca₃(PO₄)₂ : phosphate with Ca²⁺ (not alkali or NH₄⁺), insoluble.
Mg(OH)₂ : hydroxide with Mg²⁺ (not an exception), insoluble.
CaCO₃ : carbonate with Ca²⁺ (not alkali or NH₄⁺), insoluble.
The soluble ones: NaOH (alkali metal hydroxide), HCl (strong acid, fully dissociates), Na₂S (alkali metal sulfide), Pb(NO₃)₄ (nitrate, always soluble), Al(NO₃)₃ (nitrate, always soluble).
Write the formulas of the two reactants (both aqueous ionic compounds).
Identify the ions present: two cations and two anions.
Swap the partners: pair each cation with the other anion.
Write the correct formulas of the two possible products (remember to use the right ion charges to get the formula).
Check each product against the solubility rules.
If at least one product is insoluble, a precipitation reaction occurs. That product is the precipitate, written with (s). Everything else that remains dissolved gets (aq).
If both products are soluble, no net reaction occurs (write "NR" or "no reaction").
Ions present: Ag⁺, NO₃⁻, Mg²⁺, SO₄²⁻.
Swap partners: Ag⁺ pairs with SO₄²⁻ to give Ag₂SO₄. Mg²⁺ pairs with NO₃⁻ to give Mg(NO₃)₂.
Solubility check: Ag₂SO₄ is insoluble (sulfate exception for Ag⁺ is not standard, but in fact Ag₂SO₄ is only slightly soluble and is treated as insoluble in this course). Mg(NO₃)₂ is soluble (nitrates always soluble).
Balanced molecular equation: 2 AgNO₃(aq) + MgSO₄(aq) → Ag₂SO₄(s) + Mg(NO₃)₂(aq)
Precipitate: Ag₂SO₄, which is colourless/white.
Ions present: Na⁺, NO₃⁻, Mg²⁺, Cl⁻.
Swap partners: Na⁺ with Cl⁻ gives NaCl. Mg²⁺ with NO₃⁻ gives Mg(NO₃)₂.
Solubility check: NaCl is soluble (alkali metal salt). Mg(NO₃)₂ is soluble (nitrate).
Both products are soluble, so no reaction occurs. Write NR.
Ions present: Na⁺, Cl⁻, Ag⁺, NO₃⁻.
Swap partners: Na⁺ with NO₃⁻ gives NaNO₃. Ag⁺ with Cl⁻ gives AgCl.
Solubility check: NaNO₃ is soluble. AgCl is insoluble (halide with Ag⁺).
Balanced molecular equation: NaCl(aq) + AgNO₃(aq) → AgCl(s) + NaNO₃(aq)
Precipitate: AgCl, which is white. At the molecular level, the precipitate settles to the bottom of the beaker, while Na⁺ and NO₃⁻ remain as dissociated spectator ions floating in solution.
Writing the molecular equation: Reactants with (aq), precipitate with (s), gases with (g), liquids with (l).
Writing the complete ionic equation: Split every soluble ionic compound into its ions. Leave the precipitate as a complete formula.
Example for NaCl + AgNO₃: Na⁺(aq) + Cl⁻(aq) + Ag⁺(aq) + NO₃⁻(aq) → AgCl(s) + Na⁺(aq) + NO₃⁻(aq)
Writing the net ionic equation: Cancel the spectator ions (Na⁺ and NO₃⁻ appear on both sides). Ag⁺(aq) + Cl⁻(aq) → AgCl(s)
This is the reaction stripped to its essentials. Any time you mix a source of Ag⁺ with a source of Cl⁻ in water, you will get this same net ionic equation.
Molecular-level drawing (Question 12 style): The precipitate (AgCl, white solid) sits at the bottom of the beaker. The spectator ions (Na⁺ and NO₃⁻) float freely in solution above it, shown with their charges. Water molecules are omitted by convention in these diagrams.
Precipitation reactions are used in water treatment plants to remove heavy metals and other contaminants. Adding a reagent that forms an insoluble compound with the contaminant pulls it out of solution as a solid, which can then be filtered off. The same principle is used in qualitative analysis in the lab: by adding specific reagents in sequence, you can identify which ions are present in an unknown solution based on which precipitates form and what colour they are.
Students often assume that mixing any two salt solutions produces a precipitate. Many combinations yield no reaction at all, because both possible products are soluble.
Confusing "insoluble" with "does not exist." Insoluble compounds exist as solids; they simply do not dissolve in water.
Forgetting the exceptions to solubility rules. The most commonly tested exceptions: halides with Ag⁺, Pb²⁺, Hg₂²⁺ are insoluble; sulfates with Ba²⁺ and Pb²⁺ are insoluble.
Writing the precipitate with (aq) instead of (s). If it is the precipitate, it is a solid: always (s).
⚠️ Solubility rules will be tested directly (circle the insoluble compounds) and indirectly (predict whether a precipitate forms).
⚠️ You must be able to write molecular, complete ionic, and net ionic equations. All three forms may appear on the exam.
⚠️ Know common precipitate colours: AgCl is white, AgI is pale yellow, PbI₂ is bright yellow, BaSO₄ is white, CaCO₃ is white, Mg(OH)₂ is white.
⚠️ "No reaction" is a valid answer. If you determine both products are soluble, say so confidently.
True or False: All nitrate salts are soluble in water.
Fill in the blank: AgBr is ____ (soluble / insoluble) because halides with Ag⁺ are an exception.
True or False: If you mix NaCl(aq) and KNO₃(aq), a precipitate forms.
Fill in the blank: In a net ionic equation, ____ ions are removed because they appear unchanged on both sides.
True or False: BaSO₄ is soluble in water.
Answers: 1. True. 2. Insoluble. 3. False (NaNO₃ and KCl are both soluble; no reaction). 4. Spectator. 5. False (sulfate exception for Ba²⁺).
Q: You mix aqueous solutions of FeCl₃ and NaOH. Predict whether a precipitate forms, and if so, identify it.
A: Swap ions: Fe³⁺ pairs with OH⁻ to give Fe(OH)₃. Na⁺ pairs with Cl⁻ to give NaCl. Fe(OH)₃ is insoluble (hydroxide with Fe³⁺), NaCl is soluble. Precipitate: Fe(OH)₃, a rust-brown solid.
Q: Write the net ionic equation for the reaction between Pb(NO₃)₂(aq) and KI(aq).
A: Molecular: Pb(NO₃)₂(aq) + 2 KI(aq) → PbI₂(s) + 2 KNO₃(aq). Cancel K⁺ and NO₃⁻ spectator ions. Net ionic: Pb²⁺(aq) + 2 I⁻(aq) → PbI₂(s).
Q: Is Ca₃(PO₄)₂ soluble or insoluble? Explain using the solubility rules.
A: Insoluble. Phosphates are generally insoluble, and Ca²⁺ is not an alkali metal or NH₄⁺, so it does not qualify for the exceptions.
Q: A student mixes Ba(OH)₂(aq) and Na₂SO₄(aq) and observes a white solid forming. What is the precipitate? Write the balanced molecular equation.
A: The precipitate is BaSO₄ (white). Balanced equation: Ba(OH)₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2 NaOH(aq).
Solubility and precipitation connect directly to the balancing equations material in the companion study notes for this week: you cannot write a correct precipitation equation without first being able to balance it. These concepts also feed forward into stoichiometry (how much precipitate forms from a given amount of reactant), acid-base chemistry (neutralisation reactions are a special case of double displacement), and electrochemistry (where knowing which ions are in solution matters for cell potentials). In lab, gravimetric analysis relies entirely on precipitation reactions to determine the amount of an analyte.
Solubility rules, precipitation reaction, insoluble compounds, precipitate, molecular equation, complete ionic equation, net ionic equation, spectator ions, double displacement, metathesis, AgCl, BaSO₄, PbI₂, qualitative analysis, gravimetric analysis, CHM 11100, general chemistry, Purdue chemistry, aqueous solutions, ion exchange