Solubility Rules

Why This Matters

Solubility rules let you predict what happens when ionic compounds meet water. That prediction is the foundation for writing net ionic equations, identifying precipitation reactions, and working with equilibrium systems like $K_{sp}$.

On the AP exam, you'll need to look at two aqueous solutions mixing and immediately know whether a solid crashes out or everything stays dissolved. You'll also apply these rules when identifying spectator ions, writing net ionic equations, calculating molar solubility from $K_{sp}$, and predicting how common ions shift equilibria. These rules connect to reaction stoichiometry, equilibrium calculations, and acid-base chemistry (since pH affects solubility).

Don't just memorize which compounds dissolve. Understanding why certain ion combinations form insoluble precipitates while others don't is what the AP exam really tests.

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The "Always Soluble" Cations

Some cations attract anions so weakly in the solid lattice that their compounds dissolve no matter what anion they're paired with. These cations have low charge density: their +1 charge is spread over a relatively large ionic radius, so the lattice energy is easily overcome by hydration energy.

Alkali Metal Compounds

• All Group 1 cations ($Li^+$, $Na^+$, $K^+$, $Rb^+$, $Cs^+$) form soluble compounds. No exceptions on the AP exam.
• Low charge density means weak electrostatic attraction to anions, so $\Delta H_{hydration}$ easily overcomes $\Delta H_{lattice}$.
• Exam shortcut: If you see $Na^+$ or $K^+$ in a compound, it dissolves. These ions are almost always spectator ions in net ionic equations.

Ammonium Compounds

• $NH_4^+$ salts are always soluble. Treat ammonium like an alkali metal for solubility purposes.
• The polyatomic structure distributes the +1 charge across multiple atoms, creating low charge density similar to $K^+$.
• Ammonium salts often show up on exams as the soluble source of otherwise "insoluble" anions like $CO_3^{2-}$ or $PO_4^{3-}$.

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The "Always Soluble" Anions

Certain anions form such weak ionic bonds with nearly all cations that their salts dissolve universally. These anions typically have charge distributed across multiple atoms through resonance, making lattice formation unfavorable.

Nitrates

• All nitrate ($NO_3^-$) compounds are soluble. This is the single most reliable "always soluble" anion rule.
• Resonance stabilization delocalizes the negative charge across three oxygen atoms, weakening cation-anion attraction in the lattice.
• When you need a soluble source of any cation (like $Pb^{2+}$ or $Ag^+$), reach for its nitrate salt.

Acetates

• All acetate ($CH_3COO^-$ or $C_2H_3O_2^-$) compounds are soluble. The carboxylate group distributes charge effectively.
• Acetate is the conjugate base of acetic acid, which links solubility rules to acid-base equilibria.
• In acidic solutions, acetate can protonate to form acetic acid, potentially affecting precipitation reactions.

Chlorates and Perchlorates

• Chlorate ($ClO_3^-$) and perchlorate ($ClO_4^-$) salts are soluble. These are large anions with highly delocalized charge.
• These anions also appear in redox contexts as oxidizing agents, so recognizing their solubility helps you focus on the actual reaction.
• Less commonly tested than nitrates, but they follow the same "always soluble" logic.

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Halides: Soluble with Key Exceptions

Chlorides, bromides, and iodides follow a "mostly soluble" pattern, but specific cations form insoluble precipitates. The exceptions occur when cations have high polarizing power or develop significant covalent character with the halide, stabilizing the solid lattice.

Chlorides, Bromides, and Iodides (General Rule)

• Most halide salts ($Cl^-$, $Br^-$, $I^-$) are soluble, including common salts like $NaCl$, $KBr$, and $CaCl_2$.
• Soluble halides are strong electrolytes: they dissociate completely, making them useful for conductivity and colligative property problems.
• Memorize the exceptions: $Ag^+$, $Pb^{2+}$, and $Hg_2^{2+}$ form insoluble halides.

Silver Halides

• $AgCl$, $AgBr$, and $AgI$ are insoluble. These are classic precipitation products in qualitative analysis.
• $K_{sp}$ values decrease from $AgCl$ to $AgI$, meaning iodide precipitates silver most completely.
• Silver halides decompose in light (photosensitivity), which is why they were historically used in photographic film.

Lead(II) Halides

• $PbCl_2$, $PbBr_2$, and $PbI_2$ are insoluble at room temperature, though $PbCl_2$ has moderate solubility in hot water.
• Temperature dependence: $PbCl_2$ solubility increases significantly with heating. Since dissolution is endothermic here, adding heat shifts the equilibrium toward dissolved ions, a classic Le Châtelier's principle question.
• $PbI_2$ forms a distinctive bright yellow solid, commonly tested in precipitation identification.

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Sulfates: Soluble with Alkaline Earth Exceptions

Sulfate compounds are generally soluble, but several Group 2 cations and lead form insoluble precipitates. The doubly-charged sulfate anion forms strong lattice structures with doubly-charged cations of similar size, where lattice energy wins out over hydration energy.

Sulfates (General Rule)

• Most sulfate ($SO_4^{2-}$) compounds are soluble, including common salts like $Na_2SO_4$, $CuSO_4$, and $(NH_4)_2SO_4$.
• The 2- charge makes sulfate a good test case for understanding lattice energy trends.
• Memorize the exceptions: $Ba^{2+}$, $Sr^{2+}$, $Pb^{2+}$, and $Ca^{2+}$ (calcium sulfate is only slightly soluble).

Barium and Strontium Sulfates

• $BaSO_4$ and $SrSO_4$ are insoluble. Barium sulfate has an extremely low $K_{sp}$ (about $1.1 \times 10^{-10}$), making it essentially completely insoluble.
• Medical application: $BaSO_4$ is used as a contrast agent for X-rays because it's opaque to X-rays and won't dissolve in the body (so toxic $Ba^{2+}$ ions never enter the bloodstream).
• Gravimetric analysis: Precipitating $BaSO_4$ is a classic method for quantifying sulfate concentration in a sample.

Lead(II) Sulfate

• $PbSO_4$ is insoluble. Note that lead appears as an exception in multiple solubility categories (halides, sulfates, and more). If you see $Pb^{2+}$ on the exam, be suspicious of solubility.
• Battery chemistry: Lead sulfate forms during discharge of lead-acid batteries, connecting solubility to electrochemistry.
• Common ion effect: Adding excess sulfate shifts the equilibrium to precipitate more $PbSO_4$, a testable $K_{sp}$ concept.

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The "Generally Insoluble" Anions

Hydroxides, carbonates, phosphates, and sulfides form insoluble compounds with most cations. These anions either have high charge density (creating strong lattice energies) or are conjugate bases of weak acids (making their solubility pH-dependent). The only cations that keep these anions in solution are the "always soluble" ones: alkali metals and $NH_4^+$.

Hydroxides

• Most hydroxides ($OH^-$) are insoluble. The exceptions are alkali metal hydroxides and $Ba(OH)_2$.
• $NaOH$, $KOH$, and $Ba(OH)_2$ are the common soluble strong bases you need to know. $Ca(OH)_2$ and $Sr(OH)_2$ are slightly soluble.
• pH connection: Insoluble hydroxides like $Al(OH)_3$ and $Fe(OH)_3$ dissolve in acidic solutions because $H^+$ neutralizes $OH^-$, pulling the dissolution equilibrium forward.

Carbonates

• Most carbonates ($CO_3^{2-}$) are insoluble. Only alkali metal and ammonium carbonates dissolve.
• Acid sensitivity: Carbonates react with acids to produce $CO_2$ gas, so their solubility increases dramatically at low pH. The gas escaping the solution drives the reaction forward (Le Châtelier's principle).
• Common precipitates: $CaCO_3$ (limestone/chalk), $BaCO_3$, and $PbCO_3$ are frequently tested insoluble carbonates.

Phosphates

• Most phosphates ($PO_4^{3-}$) are insoluble. The high 3- charge creates very strong lattice energies with multivalent cations.
• Biological relevance: Calcium phosphate ($Ca_3(PO_4)_2$) is the main mineral component of bones and teeth.
• $K_{sp}$ calculations: The 3:2 stoichiometric ratio between $Ca^{2+}$ and $PO_4^{3-}$ makes these excellent practice for complex equilibrium problems. If $s$ is the molar solubility, then $K_{sp} = (3s)^3(2s)^2 = 108s^5$.

Sulfides

• Most sulfides ($S^{2-}$) are insoluble. Exceptions are alkali metals, ammonium, and alkaline earth metals.
• Qualitative analysis: Metal sulfides have distinctive colors ($CuS$ is black, $CdS$ is yellow, $MnS$ is pink), useful for identification.
• pH dependence: Sulfide is a weak base, so $H_2S$ equilibria affect sulfide precipitation in acidic solutions. Lowering pH reduces $S^{2-}$ concentration, which means only the least soluble metal sulfides precipitate from acidic solution.

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Quick Reference Table

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Self-Check Questions

1. Which two categories of cations always form soluble compounds regardless of the anion, and what property do they share that explains this behavior?

2. You mix solutions of $Pb(NO_3)_2$ and $NaCl$. Identify the precipitate, write the net ionic equation, and explain why nitrate and sodium are spectator ions.

3. Compare the solubility behavior of $CaSO_4$ and $BaSO_4$. How would you use this difference to selectively precipitate one cation from a mixture?

4. A student claims that all chlorides are soluble. Provide three specific counterexamples and explain what these exceptions have in common.

5. An FRQ asks how decreasing pH affects the solubility of $Fe(OH)_3$. Using Le Châtelier's principle and the concept of weak base anions, explain whether more or less solid dissolves.