Solubility Rules
Which ionic compounds dissolve in water, with the exceptions that actually get tested.
Showing 14 of 14.
| Ion or family | Formula | Usually | Exceptions |
|---|---|---|---|
| Group 1 metals | Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺ | Soluble | None |
| Ammonium | NH₄⁺ | Soluble | None |
| Nitrate | NO₃⁻ | Soluble | None |
| Acetate | C₂H₃O₂⁻ | Soluble | Ag⁺ is only slightly soluble |
| Chlorate and perchlorate | ClO₃⁻, ClO₄⁻ | Soluble | None |
| Chloride, bromide, iodide | Cl⁻, Br⁻, I⁻ | Soluble | Ag⁺, Pb²⁺, Hg₂²⁺ |
| Sulfate | SO₄²⁻ | Soluble | Ba²⁺, Pb²⁺, Sr²⁺; Ca²⁺ and Ag⁺ slightly |
| Carbonate | CO₃²⁻ | Insoluble | Group 1 and NH₄⁺ |
| Phosphate | PO₄³⁻ | Insoluble | Group 1 and NH₄⁺ |
| Chromate | CrO₄²⁻ | Insoluble | Group 1, NH₄⁺, Ca²⁺, Mg²⁺ |
| Sulfide | S²⁻ | Insoluble | Group 1, NH₄⁺, and Group 2 |
| Hydroxide | OH⁻ | Insoluble | Group 1, NH₄⁺, Ba²⁺; Ca²⁺ and Sr²⁺ slightly |
| Oxide | O²⁻ | Insoluble | Group 1; Group 2 oxides react with water |
| Sulfite | SO₃²⁻ | Insoluble | Group 1 and NH₄⁺ |
When two rules disagree, the soluble rule wins. AgNO₃ contains silver, which makes halides insoluble, but nitrate is soluble without exception — so silver nitrate dissolves. “Slightly soluble” compounds such as CaSO₄ sit between the two and are usually treated as insoluble in a precipitation question.
Solubility rules say which ionic compounds dissolve in water. Group 1 metals, ammonium and nitrates are always soluble; carbonates, phosphates and hydroxides generally are not. Mixing two solutions whose swap produces an insoluble pair gives a precipitate — the solid that appears is the whole point of the table.
What the rules are actually for
On their own the rules answer a narrow question: will this compound dissolve. Their real use is predicting what happens when two solutions are mixed. Swap the partners, check each new pairing against the table, and if either is insoluble that is your precipitate — the visible solid that forms and settles out.
That prediction is also the first step in writing a net ionic equation. Anything soluble is already floating around as separate ions and takes no part in the reaction, so it is crossed out as a spectator. What remains — the ions that actually combine into the solid — is the net ionic equation.
Solubility is a spectrum, not a switch
Nothing is completely insoluble; the labels are a convenient simplification of a continuous quantity. Compounds described as slightly soluble, such as calcium sulfate, dissolve enough to matter in some contexts and not in others, which is why different textbooks occasionally place them on opposite sides. For predicting a precipitate they are treated as insoluble; for a quantitative answer the solubility product K_sp replaces the table entirely.
- 1 Write both compounds as their ions. Mixing AgNO₃ with NaCl gives Ag⁺, NO₃⁻, Na⁺ and Cl⁻ in the same beaker.
- 2 Swap the partners. The two possible new pairings are AgCl and NaNO₃.
- 3 Check each against the table. Chlorides are soluble except with Ag⁺, so AgCl is insoluble; sodium and nitrate are both always soluble, so NaNO₃ stays dissolved.
- 4 Name the precipitate. AgCl comes out as a white solid; the sodium and nitrate ions are spectators.
- 5 Write the net ionic equation. Ag⁺(aq) + Cl⁻(aq) → AgCl(s), with the spectator ions left out.
The order to apply the rules in
Work down the list and stop at the first rule that applies — the earlier rules override the later ones.
| Check | Rule |
|---|---|
| 1. Group 1 or ammonium? | Soluble, with no exceptions |
| 2. Nitrate, chlorate, perchlorate or acetate? | Soluble |
| 3. Halide (Cl⁻, Br⁻, I⁻)? | Soluble except Ag⁺, Pb²⁺, Hg₂²⁺ |
| 4. Sulfate? | Soluble except Ba²⁺, Pb²⁺, Sr²⁺ (Ca²⁺, Ag⁺ slightly) |
| 5. Anything else | Generally insoluble — carbonate, phosphate, sulfide, hydroxide, oxide |
Where students lose marks
The biggest one is applying the rules in the wrong order. Silver nitrate looks like it should be insoluble because silver makes halides insoluble — but nitrate is soluble without exception, and the soluble rule comes first. AgNO₃ dissolves; AgCl does not. Two rules only conflict when one of them has no exceptions, and that is the one that wins.
The second is forgetting that a precipitate needs the ions to actually meet. Mixing two solutions that share an ion, or mixing two soluble compounds whose swap is also soluble, produces no reaction at all — just a mixture of ions. Writing an equation with a solid in it when nothing precipitated is a more common error than missing a precipitate that did form.