Solubility Rules: Predicting Which Ionic Compounds Dissolve in Water

Why Solubility Rules Exist

When two aqueous solutions of ionic compounds are mixed, cations and anions swap partners. If one of the new combinations is insoluble in water, a solid precipitate forms and the reaction proceeds. If all four ions stay happily dissolved, nothing visible happens.

Predicting precipitates requires knowing which ionic compounds dissolve and which do not. Chemists summarize this knowledge in a set of solubility rules — empirical guidelines that cover the overwhelming majority of the compounds you will meet in intro chemistry.

These rules are memorization-heavy but short. Once you know them, you can predict precipitation reactions in your head.

The Soluble Rules

A compound is soluble (dissolves well in water) if it contains any of the following ions.

Always soluble (with essentially no exceptions)

• Group 1 cations (Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺)
• Ammonium (NH₄⁺)
• Nitrate (NO₃⁻)
• Acetate (C₂H₃O₂⁻ or CH₃COO⁻)
• Perchlorate (ClO₄⁻)
• Chlorate (ClO₃⁻)

Any compound containing one of these ions dissolves in water. Sodium anything, potassium anything, ammonium anything, nitrate anything — soluble. This is why nitrates are so useful in lab: they keep the counter-ion in solution no matter what else is happening.

Soluble with a few exceptions

• Chloride (Cl⁻), Bromide (Br⁻), Iodide (I⁻) — soluble except with Ag⁺, Pb²⁺, Hg₂²⁺.
• Sulfate (SO₄²⁻) — soluble except with Ba²⁺, Sr²⁺, Pb²⁺, Ca²⁺ (calcium sulfate is only slightly soluble), and a few others.

The exceptions are the precipitates that drop out when you mix a chloride-containing solution with silver nitrate (AgCl), or when you mix a sulfate-containing solution with barium chloride (BaSO₄). Both reactions are used as analytical tests for the corresponding ions.

The Insoluble Rules

Compounds containing the following anions are generally insoluble, except when paired with a "always soluble" cation (Group 1 cations or ammonium).

• Carbonate (CO₃²⁻)
• Phosphate (PO₄³⁻)
• Sulfide (S²⁻) — exceptions also include Group 2 cations
• Hydroxide (OH⁻) — soluble exceptions include NaOH, KOH, and the heavier Group 2 hydroxides (Ba(OH)₂ is soluble, Ca(OH)₂ is only slightly soluble)
• Oxide (O²⁻) — most oxides are insoluble; when they do "dissolve" they usually react with water to form hydroxides

In plain English: if you see a compound that contains one of these anions paired with anything other than Na⁺, K⁺, Li⁺, or NH₄⁺, assume it is a precipitate.

A Compact Summary Table

If you want one table to refer to, this version captures the most important rules.

| Ion | Rule | Exceptions | |---|---|---| | Na⁺, K⁺, Li⁺, Rb⁺, Cs⁺, NH₄⁺ | Always soluble | none | | NO₃⁻, ClO₄⁻, ClO₃⁻, C₂H₃O₂⁻ | Always soluble | none | | Cl⁻, Br⁻, I⁻ | Soluble | AgX, PbX₂, Hg₂X₂ insoluble | | SO₄²⁻ | Soluble | BaSO₄, SrSO₄, PbSO₄, Ca²⁺ partially | | CO₃²⁻, PO₄³⁻, S²⁻ | Insoluble | Group 1 + NH₄⁺ soluble | | OH⁻ | Insoluble | NaOH, KOH, Ba(OH)₂ soluble; Ca(OH)₂ slightly |

Using the Rules to Predict a Precipitate

Example 1: Mix aqueous solutions of silver nitrate (AgNO₃) and sodium chloride (NaCl).

1. Identify possible products by swapping partners: AgCl and NaNO₃.
2. Check solubility:
   • AgCl — insoluble (silver with chloride is an exception).
   • NaNO₃ — soluble (Na⁺ is always soluble, NO₃⁻ is always soluble).
3. AgCl will precipitate; NaNO₃ stays in solution.

Balanced molecular equation: AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)

Example 2: Mix aqueous solutions of sodium nitrate (NaNO₃) and potassium chloride (KCl).

1. Possible products: NaCl and KNO₃.
2. Check solubility:
   • NaCl — soluble.
   • KNO₃ — soluble.
3. No precipitate. No reaction occurs.

Writing Net Ionic Equations

Once you have identified the precipitate, the net ionic equation strips away the spectator ions — ions that appear unchanged on both sides.

For AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq):

• Full ionic: Ag⁺(aq) + NO₃⁻(aq) + Na⁺(aq) + Cl⁻(aq) → AgCl(s) + Na⁺(aq) + NO₃⁻(aq)
• Spectators: Na⁺ and NO₃⁻ cancel.
• Net ionic: Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

The net ionic equation captures the actual chemistry — two ions coming together to form a solid — without the clutter of ions that are just watching.

What to Actually Memorize

Focus on these high-value rules first:

1. Group 1 cations and ammonium compounds are always soluble.
2. Nitrates are always soluble.
3. Silver and lead halides are insoluble. (Gives you AgCl, PbCl₂, etc.)
4. Barium sulfate is insoluble. (Common precipitation test.)
5. Most carbonates, phosphates, and hydroxides are insoluble unless paired with a Group 1 cation or ammonium.

These five rules cover the majority of intro-chemistry precipitation problems. Add the remaining exceptions as you encounter them.

The Rules Are Not Perfect

Real solubility is a continuous quantity — compounds are more or less soluble, not simply "soluble" or "insoluble." "Slightly soluble" compounds like Ca(OH)₂ and CaSO₄ sit on the boundary and behave differently in different problems. For quantitative work, chemists use the solubility product constant (K_sp), which is a topic in its own right. For the level of chemistry where you are reading solubility rules, the yes/no guidelines in this article are the right tool for the job.