Solubility rules: predicting what dissolves
Solubility rules predict which ionic compounds dissolve in water. Study the always-soluble groups, the key exceptions, and a compact reference table.
Quick answer
Solubility rules are a set of empirical guidelines that predict whether an ionic compound dissolves in water. Nitrates, acetates, and Group 1 / ammonium salts are always soluble; halides and sulfates are soluble with a few exceptions; carbonates, phosphates, hydroxides, oxides, and sulfides are insoluble except with Group 1 cations or NH4+.
Unit: qualitative (no units)
Example: NaCl is soluble (Group 1, and Cl− is a halide with no exception for Na+). AgCl is insoluble (Ag+ is a halide exception). CaCO3 is insoluble (carbonate without a Group 1 or NH4+ cation).
What you will learn
- State the solubility rules for the common ion families
- Identify the key exceptions for halides, sulfates, and hydroxides
- Use the rules to classify any common ionic compound as soluble or insoluble
- Apply the rules in the correct priority order for fast predictions
- Recognize borderline cases where Ksp refines the qualitative rules
Why solubility rules exist
When two clear solutions are mixed and a solid appears, solubility rules told you it would happen. These rules are empirical patterns — summaries of thousands of solubility measurements — that let you classify common ionic compounds as soluble or insoluble without looking anything up. They are the foundation of precipitation reactions and qualitative analysis.
The rules are organized by anion family, with the cations that break each pattern memorized as short exception lists. Learn the always-soluble groups first, then the exceptions, and you can classify nearly any common salt in seconds.
Always soluble: the no-exception groups
These ions form soluble compounds with essentially every cation — learn them as the no-exception list:
- Nitrates (NO3−): all soluble. This is why nitric acid dissolves so many samples.
- Acetates (C2H3O2−): all soluble.
- Group 1 cations (Li+, Na+, K+, and the rest): all soluble.
- Ammonium (NH4+): all soluble.
- Perchlorates (ClO4−): all soluble (often included with the always-soluble group).
If either ion of a compound falls in this list, the compound dissolves. Na2CO3 is soluble because of Na+, even though carbonates are usually insoluble.
Mostly soluble: halides and sulfates
Two anion families are soluble in general but have famous exceptions worth memorizing:
- Halides (Cl−, Br−, I−): soluble except with Ag+, Pb2+, and Hg2 2+. So NaCl dissolves but AgCl, PbBr2, and Hg2Cl2 precipitate.
- Sulfates (SO4 2−): soluble except with Ba2+, Pb2+, and Sr2+. Calcium sulfate (CaSO4) is the borderline case — slightly soluble, often grouped with the exceptions.
Note that fluoride (F−) is sometimes treated separately: many fluorides, like CaF2 and MgF2, are insoluble. When a problem says halides, check whether it means Cl−/Br−/I− or all four.
Mostly insoluble: carbonates, phosphates, hydroxides, and friends
The remaining common anions form compounds that are insoluble unless paired with a Group 1 cation or NH4+:
- Carbonates (CO3 2−): insoluble except Group 1 and NH4+. (CaCO3 — limestone — is the classic example.)
- Phosphates (PO4 3−): insoluble except Group 1 and NH4+.
- Hydroxides (OH−): insoluble except Group 1 and NH4+; Ca(OH)2, Sr(OH)2, and Ba(OH)2 are moderately soluble.
- Oxides (O2−): insoluble except Group 1 (which react with water to form hydroxides).
- Sulfides (S2−): insoluble except Group 1, NH4+, and the alkaline earths CaS, SrS, BaS.
The pattern to internalize: if the anion is not in the always-soluble or mostly-soluble lists, assume insoluble until a Group 1 or ammonium cation rescues it.
The study table
Memorize this compact study table. Apply the rows top to bottom — the first row that mentions one of your ions decides the outcome.
| Ion or family | Solubility | Exceptions |
|---|---|---|
| NO3− (nitrates) | Soluble | None |
| C2H3O2− (acetates) | Soluble | None |
| Group 1+, NH4+ | Soluble | None |
| Cl−, Br−, I− (halides) | Soluble | Ag+, Pb2+, Hg2 2+ |
| SO4 2− (sulfates) | Soluble | Ba2+, Pb2+, Sr2+ (CaSO4 marginal) |
| CO3 2− (carbonates) | Insoluble | Group 1+, NH4+ |
| PO4 3− (phosphates) | Insoluble | Group 1+, NH4+ |
| OH− (hydroxides) | Insoluble | Group 1+, NH4+ (Ca/Sr/Ba moderately) |
| S2− (sulfides) | Insoluble | Group 1+, NH4+, CaS/SrS/BaS |
Worked classifications: K2SO4 — soluble (Group 1, and sulfate has no K+ exception). PbI2 — insoluble (Pb2+ halide exception). (NH4)2CO3 — soluble (NH4+ rescues the carbonate). Fe(OH)3 — insoluble. BaSO4 — insoluble (Ba2+ sulfate exception).
Limits and refinements
The rules predict the common cases, but three refinements matter. First, insoluble really means slightly soluble: every precipitate maintains a tiny equilibrium concentration of ions, which the solubility product constant Ksp measures. Second, temperature shifts borderline cases — hot water dissolves most solids better, so a slightly soluble salt may precipitate on cooling.
Third, the common-ion effect changes real solubilities: AgCl is less soluble in NaCl solution than in pure water because the extra Cl− pushes the dissolution equilibrium backward. For qualitative predictions the rules suffice; for numbers, reach for Ksp.
Common mistakes
'Insoluble' in these rules means very slightly soluble — typically less than about 0.01 mol/L. Even classic precipitates like AgCl dissolve to a tiny extent, which is exactly what the solubility product Ksp quantifies.
This one is essentially true for common salts — Group 1 cations and NH4+ have no exceptions in the standard rules. It is the safest rule on the list and the reason sodium and potassium salts are used when you need a guaranteed-soluble counterion.
Insoluble compounds still establish an equilibrium with a small concentration of dissolved ions. That is why precipitates can re-dissolve when conditions change and why Ksp expressions are written for 'insoluble' salts.
They describe the same reality at different resolutions. A compound the rules call insoluble will have a very small Ksp; a compound the rules call soluble will have a large one. Borderline cases like CaSO4 sit near the boundary and are where Ksp adds real value.
Real-world applications
Predicting whether mixing two solutions will form a precipitate
Writing net ionic equations by identifying the insoluble product
Qualitative analysis schemes that separate metal ions by selective precipitation
Water treatment: precipitating dissolved metals and phosphates for removal
Geology: explaining mineral deposit formation from groundwater
Medicine: understanding kidney stone formation as a precipitation process
Frequently asked questions
Which salts are always soluble?
Nitrates and acetates are always soluble, with no exceptions. That is why nitric acid and acetic acid are used to dissolve samples in qualitative analysis — their salts stay in solution. Every common nitrate (NaNO3, AgNO3, Pb(NO3)2, and so on) and every common acetate dissolves in water.
What are the exceptions for halides and sulfates?
Halides (Cl−, Br−, I−) are generally soluble, but three cations break the rule: Ag+, Pb2+, and Hg2 2+. So NaCl dissolves, but AgCl, PbCl2, and Hg2Cl2 precipitate. Sulfates follow a similar pattern with Ba2+, Pb2+, and Sr2+ as the main exceptions, plus CaSO4 which is only slightly soluble.
Are all Group 2 hydroxides really insoluble?
It is a slight oversimplification. Ca(OH)2, Sr(OH)2, and Ba(OH)2 are moderately soluble — Ba(OH)2 dissolves well enough to make a strongly basic solution. Mg(OH)2 is the classic nearly-insoluble example (milk of magnesia is a suspension). The rule 'hydroxides are insoluble except Group 1 and NH4+' holds if you read 'soluble' as 'very soluble.'
What is the fastest way to apply the rules on an exam?
Use them in this order: (1) Group 1 and NH4+ salts are soluble, no exceptions. (2) Nitrates, acetates, and perchlorates are soluble, no exceptions. (3) Halides are soluble except Ag+, Pb2+, Hg2 2+. (4) Sulfates are soluble except Ba2+, Pb2+, Sr2+ (CaSO4 marginal). (5) Everything else common — carbonates, phosphates, sulfides, oxides, hydroxides — is insoluble except with Group 1 or NH4+.
Do the rules ever fail?
Solubility rules predict the common cases correctly but miss borderline compounds like CaSO4, Mg(OH)2, and temperature-dependent solubilities. For quantitative work chemists use the solubility product constant Ksp, which gives actual molar solubilities. Think of the rules as a fast qualitative screen and Ksp as the precise follow-up.
How do temperature and pressure affect these rules?
Temperature matters most for borderline cases: most solids dissolve better in hot water, so a slightly soluble salt may stay dissolved when heated and precipitate on cooling. Pressure has little effect on solid solubility. Concentration and the common-ion effect also shift real solubilities, which is why Ksp calculations refine rule-based predictions.