Solubility Rules to Know for AP Chemistry

What You Need to Know

Solubility rules are quick patterns that let you predict whether an ionic compound dissolves (aq) or forms a precipitate (s) in water—without doing a full $K_{sp}$ calculation.

You use them constantly for:

Predicting precipitates in double-displacement (metathesis) reactions
Writing complete ionic and net ionic equations (AP loves this)
Deciding when to use $K_{sp}$ / comparing $Q$ vs $K_{sp}$

Core idea:

If a product is insoluble, it appears as (s) and “drives” the reaction.
If all products are soluble, no reaction (NR) is usually written (for precipitation problems).

AP Exam mindset: Solubility rules are “fast triage.” If you’re unsure, identify the anion rule first (like $\text{Cl}^-$ , $\text{SO}_4^{2-}$ , $\text{CO}_3^{2-}$ ), then check the cation exceptions.

Step-by-Step Breakdown

How to predict a precipitate (and write the net ionic)

Split reactants into ions (assume the given reactants are soluble unless told otherwise).
Swap partners to form the two possible products (double displacement):
- $AB + CD \rightarrow AD + CB$
Check solubility of each product using the rules + exceptions.
- Any insoluble product is a precipitate (s).
Write the balanced molecular equation with states.
Write the complete ionic equation: split aqueous strong electrolytes into ions; keep (s), (l), (g) intact.
Cancel spectator ions to get the net ionic equation.

Mini worked walkthrough

Mix $\text{BaCl}_2(aq)$ and $\text{Na}_2\text{SO}_4(aq)$ .

Ions: $\text{Ba}^{2+}, \text{Cl}^-$ and $\text{Na}^+, \text{SO}_4^{2-}$
Swap: $\text{BaSO}_4$ and $\text{NaCl}$
Solubility:
- $\text{BaSO}_4$ : sulfate exception → insoluble → precipitate
- $\text{NaCl}$ : Group 1 salt → soluble
Molecular:
- $\text{BaCl}_2(aq) + \text{Na}_2\text{SO}_4(aq) \rightarrow \text{BaSO}_4(s) + 2\,\text{NaCl}(aq)$
Net ionic:
- $\text{Ba}^{2+}(aq) + \text{SO}_4^{2-}(aq) \rightarrow \text{BaSO}_4(s)$

Decision point: If no insoluble product forms, write NR for precipitation. (A reaction might still occur as acid–base or redox, but that’s a different prompt.)

Key Formulas, Rules & Facts

The “always soluble” (memorize these first)

Rule	Compounds that are soluble	Notes / exceptions to watch
Group 1 (alkali metals)	$\text{Li}^+$ , $\text{Na}^+$ , $\text{K}^+$ , $\text{Rb}^+$ , $\text{Cs}^+$ salts	Always soluble in AP contexts
Ammonium	$\text{NH}_4^+$ salts	Always soluble
Nitrates	$\text{NO}_3^-$ salts	Always soluble
Acetates (and related)	$\text{C}_2\text{H}_3\text{O}_2^-$ (or $\text{CH}_3\text{COO}^-$ ) salts	Treated as soluble on AP (some are only moderately soluble in reality)
Perchlorates / chlorates	$\text{ClO}_4^-$ and $\text{ClO}_3^-$ salts	Common “always soluble” add-ons

Halides: usually soluble, with a short exception list

Anion family	Usually soluble as…	Key insoluble exceptions (form precipitates)
$\text{Cl}^-$ , $\text{Br}^-$ , $\text{I}^-$	Most chlorides/bromides/iodides	$\text{Ag}^+$ , $\text{Pb}^{2+}$ , $\text{Hg}_2^{2+}$ (mercury(I), “mercurous”)

Notes that show up on AP:

$\text{AgCl}$ , $\text{AgBr}$ , $\text{AgI}$ are classic precipitates.
$\text{PbCl}_2$ is only slightly soluble, but AP typically treats it as a precipitate under typical mixing conditions.

Sulfates: usually soluble, but the exceptions matter

Anion	Usually soluble as…	Key insoluble / low-solubility exceptions
$\text{SO}_4^{2-}$	Most sulfates	$\text{Ba}^{2+}$ , $\text{Sr}^{2+}$ , $\text{Pb}^{2+}$ (and often $\text{Ca}^{2+}$ as slightly soluble)

Practical AP takeaway:

If you see $\text{Ba}^{2+}$ + $\text{SO}_4^{2-}$ → you should immediately think $\text{BaSO}_4(s)$ .

“Usually insoluble” polyatomic anions (unless Group 1 or ammonium)

These are the precipitation workhorses.

Anion (usually insoluble)	Soluble when paired with…	Common precipitates
$\text{CO}_3^{2-}$ (carbonate)	Group 1 or $\text{NH}_4^+$	$\text{CaCO}_3(s)$ , $\text{BaCO}_3(s)$ , $\text{FeCO}_3(s)$
$\text{PO}_4^{3-}$ (phosphate)	Group 1 or $\text{NH}_4^+$	$\text{Ca}_3(\text{PO}_4)_2(s)$
$\text{CrO}_4^{2-}$ (chromate)	Group 1 or $\text{NH}_4^+$	$\text{BaCrO}_4(s)$ , $\text{PbCrO}_4(s)$
$\text{C}_2\text{O}_4^{2-}$ (oxalate)	Group 1 or $\text{NH}_4^+$	$\text{CaC}_2\text{O}_4(s)$

Fast rule: If it’s carbonate/phosphate/chromate/oxalate and the cation is NOT Group 1 or $\text{NH}_4^+$ → assume precipitate.

Hydroxides and sulfides: mostly insoluble (with key soluble-ish exceptions)

Anion	General rule	Exceptions you should know
$\text{OH}^-$	Insoluble	Soluble with Group 1 and $\text{NH}_4^+$ ; $\text{Ca}^{2+}$ , $\text{Sr}^{2+}$ , $\text{Ba}^{2+}$ hydroxides are more soluble (often treated as slightly to moderately soluble)
$\text{S}^{2-}$	Insoluble	Soluble with Group 1, Group 2 (especially $\text{Ca}^{2+}$ , $\text{Sr}^{2+}$ , $\text{Ba}^{2+}$ ), and $\text{NH}_4^+$ (rule-of-thumb set used in many AP courses)

How AP tends to treat these:

Hydroxides of Group 1 are clearly soluble (e.g., $\text{NaOH}$ ).
$\text{Ba(OH)}_2$ and $\text{Sr(OH)}_2$ are often treated as soluble enough to count as strong bases.
Most transition-metal hydroxides and sulfides precipitate.

Connecting to equilibrium: when solubility rules vs $K_{sp}$

Solubility rules give you a qualitative call: “likely (s)” vs “(aq).”
$K_{sp}$ gives you a quantitative call.

Relationship	Meaning
$Q < K_{sp}$	Unsaturated → no precipitate
$Q = K_{sp}$	Saturated at equilibrium
$Q > K_{sp}$	Supersaturated → precipitate forms until $Q = K_{sp}$

Examples & Applications

Example 1: Classic halide precipitate

Mix $\text{AgNO}_3(aq)$ and $\text{NaCl}(aq)$ .

Products: $\text{AgCl}$ and $\text{NaNO}_3$
$\text{AgCl}$ is insoluble (halide exception: $\text{Ag}^+$ )

Molecular:

$\text{AgNO}_3(aq) + \text{NaCl}(aq) \rightarrow \text{AgCl}(s) + \text{NaNO}_3(aq)$
Net ionic:
$\text{Ag}^+(aq) + \text{Cl}^-(aq) \rightarrow \text{AgCl}(s)$

AP twist: If asked for spectators: $\text{Na}^+$ and $\text{NO}_3^-$ .

Example 2: Sulfate exception (high-yield)

Mix $\text{K}_2\text{SO}_4(aq)$ and $\text{Ba(NO}_3)_2(aq)$ .

Products: $\text{BaSO}_4$ and $\text{KNO}_3$
$\text{BaSO}_4$ is insoluble (sulfate exception)

Net ionic:

$\text{Ba}^{2+}(aq) + \text{SO}_4^{2-}(aq) \rightarrow \text{BaSO}_4(s)$

Key insight: Nitrates and Group 1 are “spectator-ion magnets” because they almost always stay aqueous.

Example 3: Carbonate precipitate

Mix $\text{CaCl}_2(aq)$ and $\text{Na}_2\text{CO}_3(aq)$ .

Products: $\text{CaCO}_3$ and $\text{NaCl}$
Carbonates are insoluble unless Group 1 or $\text{NH}_4^+$ → $\text{CaCO}_3(s)$ forms

Net ionic:

$\text{Ca}^{2+}(aq) + \text{CO}_3^{2-}(aq) \rightarrow \text{CaCO}_3(s)$

Exam variation: They may ask which reagent removes $\text{Ca}^{2+}$ from hard water—carbonate works by precipitating $\text{CaCO}_3$ .

Example 4: “No precipitate” (NR)

Mix $\text{Na}_2\text{SO}_4(aq)$ and $\text{KCl}(aq)$ .

Possible products: $\text{NaCl}$ and $\text{K}_2\text{SO}_4$
Both are soluble (Group 1 salts)

Result: NR for a precipitation question.

Trap reminder: Don’t force a reaction just because it’s double displacement. If everything stays aqueous, there’s no precipitate-driven change.

Common Mistakes & Traps

Forgetting the “always soluble” ions (Group 1, $\text{NH}_4^+$ , $\text{NO}_3^-$ )
- What goes wrong: You predict a precipitate like $\text{Na}_2\text{CO}_3(s)$ or $\text{KCl}(s)$ .
- Why wrong: Group 1 and $\text{NH}_4^+$ salts are essentially always soluble in AP problems.
- Fix: When you see Group 1 or $\text{NH}_4^+$ in a product, default to (aq).
Mixing up sulfate vs sulfide
- What goes wrong: Treat $\text{SO}_4^{2-}$ like $\text{S}^{2-}$ (or vice versa).
- Why wrong: Sulfates are mostly soluble with a few big exceptions; sulfides are mostly insoluble with a few big exceptions.
- Fix: Say it out loud: “sulfate = usually soluble; sulfide = usually insoluble.”
Not memorizing the key halide exceptions
- What goes wrong: You call $\text{AgCl}$ soluble.
- Why wrong: $\text{Ag}^+$ , $\text{Pb}^{2+}$ , $\text{Hg}_2^{2+}$ make halides precipitate.
- Fix: Drill the exception trio until it’s automatic.
Overthinking “slightly soluble” on AP
- What goes wrong: You hesitate to write a precipitate because something is only slightly soluble (e.g., $\text{PbCl}_2$ or $\text{CaSO}_4$ ).
- Why wrong: In typical mixing scenarios, slightly soluble compounds still commonly precipitate.
- Fix: If it’s on the exception list, treat it as (s) unless the problem gives specific concentrations/ $K_{sp}$ .
Canceling the wrong ions in net ionic equations
- What goes wrong: You cancel ions that actually form the precipitate, or you forget to split strong electrolytes.
- Why wrong: Spectators are ions present unchanged on both sides.
- Fix: Only cancel identical aqueous ions on both sides; never split solids.
Assuming a precipitate forms because two solutions are mixed
- What goes wrong: You always write products even if both are soluble.
- Why wrong: Precipitation requires an insoluble product (or $Q > K_{sp}$ ).
- Fix: Check both possible products—if both are soluble, write NR.
Missing polyatomic-ion spelling/charge leads to wrong formula and wrong solubility
- What goes wrong: You write $\text{CaPO}_4$ instead of $\text{Ca}_3(\text{PO}_4)_2$ .
- Why wrong: Wrong formula can change charge balance and what you think precipitates.
- Fix: Write ions with charges first, then cross/balance to a neutral formula.
Ignoring that solubility rules assume aqueous water, not extreme conditions
- What goes wrong: You apply rules blindly in strongly acidic/basic conditions.
- Why wrong: Some “insoluble” salts (like carbonates) can be consumed by acid in broader reaction types.
- Fix: If the question is specifically “precipitation reaction,” stick to solubility rules; if acids/bases are emphasized, consider acid–base chemistry too.

Memory Aids & Quick Tricks

Trick / mnemonic	What it helps you remember	When to use it
“Group 1 and $\text{NH}_4^+$ = always soluble”	The #1 shortcut in precipitation	Any time you’re unsure
Nitrates are always soluble	$\text{NO}_3^-$ never precipitates in AP problems	Quickly identify spectators
“PMS” = $\text{Pb}^{2+}$ , $\text{Hg}_2^{2+}$ , $\text{Ag}^+$	Halide exceptions (chloride/bromide/iodide precipitates)	When you see $\text{Cl}^-$ , $\text{Br}^-$ , $\text{I}^-$
“CASTRO BEAR” (common classroom mnemonic)	Sulfate exceptions: $\text{Ca}^{2+}$ , $\text{Sr}^{2+}$ , $\text{Ba}^{2+}$ , and often $\text{Pb}^{2+}$	When checking $\text{SO}_4^{2-}$
“CO PO CROX”	$\text{CO}_3^{2-}$ , $\text{PO}_4^{3-}$ , $\text{CrO}_4^{2-}$ , $\text{C}_2\text{O}_4^{2-}$ are insoluble except Group 1/ $\text{NH}_4^+$	Fast precipitate prediction
“Hydroxides hate dissolving”	$\text{OH}^-$ usually insoluble except Group 1 and heavier Group 2	Metal hydroxide precipitates
If it contains $\text{Na}^+$ or $\text{K}^+$ , assume (aq)	Quick scan of products	When time is tight

Warning: Mnemonics vary by teacher. If your class list includes or excludes $\text{CaSO}_4$ / $\text{Ag}_2\text{SO}_4$ as exceptions, follow your course sheet—but the big ones (Ba, Sr, Pb) are the must-knows.

Quick Review Checklist

You can instantly label (aq) for anything with Group 1, $\text{NH}_4^+$ , $\text{NO}_3^-$ , **acetate**, $\text{ClO}_4^-$ , $\text{ClO}_3^-$ .
You know halides $\text{Cl}^-$ / $\text{Br}^-$ / $\text{I}^-$ are soluble **except** with $\text{Ag}^+$ , $\text{Pb}^{2+}$ , $\text{Hg}_2^{2+}$ .
You know sulfates are soluble except with $\text{Ba}^{2+}$ , $\text{Sr}^{2+}$ , $\text{Pb}^{2+}$ (and $\text{Ca}^{2+}$ often flagged as slightly soluble).
You treat carbonates, phosphates, chromates, oxalates as insoluble unless Group 1 or $\text{NH}_4^+$ .
You treat hydroxides and sulfides as insoluble unless Group 1 (and remember key Group 2 / $\text{NH}_4^+$ exceptions).
You can go from molecular → complete ionic → net ionic cleanly (cancel only spectators).
If all possible products are soluble, you confidently write NR (for precipitation).

You’ve got this—solubility rules are pure pattern recognition, and patterns are easy points when you practice them a few times.