Solubility Rules for Ionic Compounds

Why This Matters

Solubility rules are the foundation for predicting what happens when ionic compounds meet water—and that prediction skill is exactly what you're being tested on. Whether you're writing net ionic equations, identifying precipitates in double displacement reactions, or explaining why certain compounds form in qualitative analysis, these rules are your roadmap. They connect directly to precipitation reactions, equilibrium concepts, and analytical techniques that appear throughout inorganic chemistry.

Here's the key insight: solubility isn't random. It follows patterns based on ion charge density, lattice energy, and hydration energy. The cations and anions that form soluble compounds do so because water molecules can effectively stabilize those ions in solution. Don't just memorize a list—understand that each rule reflects a competition between the energy holding the crystal together and the energy released when water surrounds each ion. That conceptual framework will serve you far better on exams than rote memorization.

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Always Soluble: The Reliable Ions

These ions form soluble compounds with essentially no exceptions. Their low charge density and favorable hydration energies mean water wins the tug-of-war against lattice forces every time.

Alkali Metal Compounds ($Li^+$, $Na^+$, $K^+$, $Rb^+$, $Cs^+$)

• No exceptions exist—every alkali metal salt dissolves in water, making this your most reliable rule
• Low charge density of +1 cations means weak lattice energies that water easily overcomes
• Universal applicability means if you see $Na^+$ or $K^+$ in a compound, mark it soluble immediately

Ammonium Compounds ($NH_4^+$)

• Polyatomic cation with +1 charge behaves like alkali metals for solubility purposes
• Common in fertilizers and laboratory reagents precisely because of reliable water solubility
• Exam shortcut—ammonium salts of "insoluble" anions (carbonates, phosphates, sulfides) are still soluble

Nitrate Compounds ($NO_3^-$)

• No common exceptions—nitrate salts dissolve regardless of the cation present
• Resonance-stabilized anion with delocalized charge makes it easily hydrated
• Practical significance in fertilizers ($NH_4NO_3$) and oxidizers relies on this high solubility

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Generally Soluble with Key Exceptions: The Halides and Sulfates

These anions form soluble compounds most of the time, but specific cations create insoluble precipitates. The exceptions involve cations with high charge density or polarizing power that create strong lattice interactions.

Chlorides, Bromides, and Iodides ($Cl^-$, $Br^-$, $I^-$)

• Soluble except with $Ag^+$, $Pb^{2+}$, and $Hg_2^{2+}$—memorize these three exception cations
• Silver halides ($AgCl$, $AgBr$, $AgI$) are classic precipitates used in qualitative analysis and photography
• Lead(II) chloride ($PbCl_2$) is slightly soluble in hot water—a detail that occasionally appears on exams

Sulfate Compounds ($SO_4^{2-}$)

• More exceptions than halides—insoluble with $Ba^{2+}$, $Sr^{2+}$, $Ca^{2+}$, $Pb^{2+}$, $Ag^+$, and $Hg_2^{2+}$
• Barium sulfate ($BaSO_4$) is extremely insoluble and used in medical imaging (barium swallows)
• Calcium sulfate ($CaSO_4$) is only slightly insoluble—sometimes listed as "sparingly soluble"

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Generally Insoluble with Key Exceptions: Hydroxides, Carbonates, Phosphates, and Sulfides

These anions form insoluble compounds by default. Their higher charge densities create strong lattice energies that water typically cannot overcome—except when paired with those reliable +1 cations.

Hydroxide Compounds ($OH^-$)

• Insoluble except with alkali metals and $Ba^{2+}$—most metal hydroxides precipitate from solution
• Strong bases like $NaOH$ and $KOH$ are soluble; weak bases like $Mg(OH)_2$ and $Fe(OH)_3$ are not
• $Ca(OH)_2$ (lime water) is slightly soluble—enough to make basic solutions but not highly concentrated ones

Carbonate Compounds ($CO_3^{2-}$)

• Insoluble except with alkali metals and $NH_4^+$—the -2 charge creates strong lattice interactions
• Calcium carbonate ($CaCO_iteiteite_3$) as limestone/chalk is a classic example of carbonate insolubility
• Decomposition reactions of insoluble carbonates release $CO_2$ gas when treated with acid

Phosphate Compounds ($PO_4^{3-}$)

• Insoluble except with alkali metals and $NH_4^+$—the -3 charge makes these among the least soluble salts
• Calcium phosphate ($Ca_3(PO_4)_2$) in bones and teeth relies on this insolubility for structural stability
• Fertilizer chemistry often involves converting insoluble phosphates to soluble forms plants can absorb

Sulfide Compounds ($S^{2-}$)

• Insoluble except with alkali metals, $NH_4^+$, and alkaline earth metals—note the extra exceptions here
• Qualitative analysis exploits different sulfide solubilities to separate metal cations in solution
• Distinctive colors of metal sulfides (black $CuS$, yellow $CdS$) aid in identification

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Quick Reference Table

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Self-Check Questions

1. Which two anions form soluble compounds with every cation, including those that typically cause precipitation (like $Ag^+$ and $Ba^{2+}$)?

2. You mix solutions of $AgNO_3$ and $NaCl$. Identify the precipitate and explain which solubility rule predicts its formation.

3. Compare and contrast the solubility behavior of hydroxides versus carbonates—which has more soluble exceptions, and why might that be?

4. A student claims that all sulfates are soluble. Identify three specific compounds that disprove this claim and explain what these exception cations have in common.

5. If you needed to precipitate $Ba^{2+}$ ions from solution, would you add $Na_2SO_4$ or $NaNO_3$? Justify your answer using solubility rules, and write the net ionic equation for any reaction that occurs.