Solubility Rules
Solubility rules are a set of guidelines that predict whether an ionic compound dissolves in water (soluble) or stays as a solid (insoluble). They let you predict the products of precipitation reactions and write net ionic equations — without memorising every compound.
Rules apply to compounds in water at room temperature. "Slightly soluble" compounds are treated as insoluble for precipitation.
The solubility rules
Read each row as "compounds containing this ion are soluble / insoluble, except for the ions listed." Work down the list — the first rule that applies wins.
Compounds containing… | Solubility in water | Key exceptions |
|---|---|---|
Group 1 metals (Li⁺, Na⁺, K⁺…) and ammonium (NH₄⁺) | ✅ Soluble | none |
Nitrates (NO₃⁻), acetates (CH₃COO⁻) | ✅ Soluble | none |
Chlorides (Cl⁻), bromides (Br⁻), iodides (I⁻) | ✅ Soluble | Ag⁺, Pb²⁺, Hg₂²⁺ (insoluble) |
Sulfates (SO₄²⁻) | ✅ Soluble | Ba²⁺, Pb²⁺, Sr²⁺ (and Ca²⁺ slightly) |
Hydroxides (OH⁻) | ❌ Insoluble | Group 1 & Ba²⁺ soluble (Ca²⁺, Sr²⁺ slightly) |
Carbonates (CO₃²⁻), phosphates (PO₄³⁻), sulfides (S²⁻) | ❌ Insoluble | Group 1 & NH₄⁺ soluble |
The top four rules cover "soluble unless…"; the bottom two cover "insoluble unless…". Group 1 and ammonium salts are always soluble — that rule beats everything below it.
How to use them: predicting a precipitate
When two solutions mix, imagine swapping the partners to get two possible new compounds.
Check each new compound against the rules. If either one is insoluble, it forms a precipitate (a solid).
If both possible products are soluble, no reaction happens — all ions stay dissolved.
Worked example
Solutions of silver nitrate (AgNO₃) and sodium chloride (NaCl) are mixed. Predict what happens and write the net ionic equation.
Swap the partners: the possible products are silver chloride (AgCl) and sodium nitrate (NaNO₃).
Check NaNO₃: it contains Na⁺ (Group 1) and NO₃⁻ (nitrate) — both rules say soluble, so it stays dissolved.
Check AgCl: chlorides are soluble except with Ag⁺ — so AgCl is insoluble and forms a white precipitate.
Write the net ionic equation using only the ions that actually react (drop the spectator ions Na⁺ and NO₃⁻).
A white precipitate of AgCl forms: Ag⁺(aq) + Cl⁻(aq) → AgCl(s)
the ions left out of the net ionic equation are the spectator ions — they were soluble on both sides and never changed.
Common mistakes
Treating "slightly soluble" compounds (like CaSO₄, Ca(OH)₂) as fully soluble — in precipitation reactions they count as insoluble.
Forgetting that Group 1 and NH₄⁺ salts are always soluble — that rule overrides the carbonate/hydroxide/sulfide "insoluble" rules.
Leaving spectator ions in the net ionic equation — they belong only in the total ionic equation.
Frequently asked questions
Are all nitrates soluble?
Yes — every nitrate (NO₃⁻) salt is soluble in water, with no common exceptions. The same is true of Group 1 (alkali metal) and ammonium salts. These are the most reliable rules.
What does "insoluble" really mean here?
No ionic compound is truly 100% insoluble, but "insoluble" means so little dissolves that a visible solid (precipitate) forms. "Slightly soluble" compounds are treated the same way for predicting precipitates.
How do solubility rules relate to Ksp?
Solubility rules are the qualitative shortcut ("soluble or not?"). The solubility product Ksp is the quantitative version used at A-Level/AP: a small Ksp means very insoluble, so the rules and Ksp agree.
This concept in your exam
Open your board's guide for exam-specific technique and past-paper practice:
Predict products and write net ionic equations — step-by-step, free to try.
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