Solubility Rules

Why This Matters

Solubility rules are your gateway to predicting what happens when ionic compounds meet water—and that's the foundation for writing net ionic equations, predicting precipitation reactions, and understanding equilibrium systems like $K_{sp}$. You're being tested on your ability to look at two aqueous solutions mixing and immediately know whether a solid will crash out of solution or whether everything stays dissolved. This connects directly to reaction stoichiometry, equilibrium calculations, and even acid-base chemistry when pH affects solubility.

Here's the thing: the AP exam won't just ask you to recite rules. You'll need to apply them in context—identifying spectator ions, writing net ionic equations, calculating molar solubility from $K_{sp}$, and predicting how common ions shift equilibria. Don't just memorize which compounds dissolve; understand why certain ion combinations form insoluble precipitates while others don't. That conceptual understanding is what separates a 3 from a 5.

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The "Always Soluble" Cations

Some cations are so weakly attracted to anions in the solid lattice that their compounds dissolve regardless of what anion they're paired with. These cations have low charge density, meaning their ionic charge is spread over a relatively large radius, resulting in weak lattice energies that are easily overcome by hydration energy.

Alkali Metal Compounds

• All Group 1 cations ($Li^+$, $Na^+$, $K^+$, $Rb^+$, $Cs^+$) form soluble compounds—no exceptions to memorize here
• 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—making these ions common spectator ions in net ionic equations

Ammonium Compounds

• $NH_4^+$ salts are always soluble—treat ammonium like an alkali metal for solubility purposes
• Polyatomic structure distributes the +1 charge across multiple atoms, creating low charge density similar to $K^+$
• Common exam appearance: ammonium salts often show up 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 (resonance structures) or have low charge density, making lattice formation unfavorable.

Nitrates

• All nitrate ($NO_3^-$) compounds are soluble—this is the most reliable "always soluble" anion rule
• Resonance stabilization delocalizes the negative charge across three oxygen atoms, weakening cation-anion attraction
• Lab application: 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
• Weak acid connection: acetate is the conjugate base of acetic acid, linking solubility rules to acid-base equilibria
• pH sensitivity: 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—large anions with highly delocalized charge
• Oxidizing agents: these anions appear in redox contexts, so recognizing their solubility helps you focus on the actual reaction
• Less commonly tested than nitrates but 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 form covalent character with the halide, stabilizing the solid lattice.

Chlorides, Bromides, and Iodides (General Rule)

• Most halide salts ($Cl^-$, $Br^-$, $I^-$) are soluble—this includes common salts like $NaCl$, $KBr$, and $CaCl_2$
• Strong electrolytes: soluble halides 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—classic precipitation products in qualitative analysis
• $K_{sp}$ values decrease from $AgCl$ to $AgI$, meaning iodide precipitates silver most completely
• Photosensitivity: silver halides decompose in light, which is why they appear in photography-related contexts

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, a potential equilibrium/Le Châtelier question
• Yellow precipitate: $PbI_2$ 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 larger, doubly-charged sulfate anion forms strong lattice structures with doubly-charged cations of similar size.

Sulfates (General Rule)

• Most sulfate ($SO_4^{2-}$) compounds are soluble—including common salts like $Na_2SO_4$, $CuSO_4$, and $(NH_4)_2SO_4$
• Doubly-charged anion: 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 is essentially completely insoluble with extremely low $K_{sp}$
• Medical application: $BaSO_4$ is used as a contrast agent for X-rays because it's opaque and won't dissolve in the body
• Gravimetric analysis: precipitating $BaSO_4$ is a classic method for quantifying sulfate concentration

Lead(II) Sulfate

• $PbSO_4$ is insoluble—note that lead appears as an exception in multiple solubility categories
• Battery chemistry: lead sulfate forms during discharge of lead-acid batteries, connecting solubility to electrochemistry
• Common ion effect: adding excess sulfate shifts 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).

Hydroxides

• Most hydroxides ($OH^-$) are insoluble—exceptions are alkali metal hydroxides and $Ba(OH)_2$
• Strong bases: $NaOH$, $KOH$, and $Ba(OH)_2$ are the common soluble strong bases you'll encounter
• pH connection: insoluble hydroxides like $Al(OH)_3$ and $Fe(OH)_3$ dissolve in acidic solutions as $H^+$ neutralizes $OH^-$

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
• Common precipitates: $CaCO_3$ (limestone), $BaCO_3$, 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
• Biological relevance: calcium phosphate ($Ca_3(PO_4)_2$) is the main component of bones and teeth
• $K_{sp}$ calculations: phosphate stoichiometry (3:2 ratio with $Ca^{2+}$) makes these excellent practice for complex equilibrium problems

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), useful for identification
• pH dependence: sulfide is a weak base, so $H_2S$ equilibria affect sulfide precipitation in acidic solutions

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Quick Reference Table

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Self-Check Questions

1. Which two 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 and contrast 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 the qualitative effect and predict whether more or less solid dissolves.