Solubility rules are a set of empirical guidelines used to predict whether an ionic compound will dissolve in water. They are based on observed patterns: compounds containing certain ions (like nitrates, alkali metals, and ammonium) are almost always soluble, while others (like most carbonates, phosphates, and hydroxides) tend to be insoluble. These rules are essential for predicting precipitation reactions, understanding double displacement reactions, and performing qualitative analysis in chemistry labs.

When two aqueous ionic solutions are mixed, the ions can recombine. If any possible combination of cation and anion forms an insoluble compound (based on solubility rules), a precipitate will form. For example, mixing AgNO₃ (soluble) and NaCl (soluble) yields AgCl, which is insoluble per the halide exception rule — so a white precipitate of AgCl forms. Use our Ion Pair Solubility Checker above to test any cation-anion combination instantly.

No-exception always-soluble ions:

• Nitrate (NO₃⁻) — all nitrate salts dissolve in water
• Acetate (CH₃COO⁻) — all acetate salts are soluble
• Chlorate (ClO₃⁻) & Perchlorate (ClO₄⁻) — all are soluble
• Alkali metal ions (Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺) — all their salts dissolve
• Ammonium (NH₄⁺) — all ammonium salts are soluble

These ions have zero exceptions — any compound containing at least one of these will be soluble.

The critical exceptions every chemistry student must know:

• Halides (Cl⁻, Br⁻, I⁻): Ag⁺, Pb²⁺, Hg₂²⁺ form insoluble salts
• Sulfates (SO₄²⁻): Ca²⁺, Sr²⁺, Ba²⁺, Pb²⁺, Hg₂²⁺ form insoluble/lightly soluble sulfates
• Carbonates (CO₃²⁻) & Phosphates (PO₄³⁻): Only Group 1 & NH₄⁺ salts are soluble
• Hydroxides (OH⁻): Only Group 1, NH₄⁺, Sr²⁺, Ba²⁺ are soluble (Ca(OH)₂ slightly soluble)
• Sulfides (S²⁻): Only Group 1, NH₄⁺, Ca²⁺, Sr²⁺, Ba²⁺ are soluble

Mnemonic: "PMS — Pb²⁺, Mercury(I), Silver cause trouble for halides & sulfates!"

Use mnemonics! The most popular one is NAG SAG:

• Nitrates — Always soluble
• Acetates — Always soluble
• Group 1 & ammonium — Always soluble
• Sulfates — Mostly soluble (watch for CaSrBaPbHg₂)
• Ammonium — All salts soluble
• Group 17 halides — Mostly soluble (watch for AgPbHg₂)

Also: PMS = Pb²⁺, Mercury(I) Hg₂²⁺, Silver Ag⁺ — these three are the frequent exceptions!

• Soluble: Dissolves readily in water at room temperature; typically >1g per 100mL water. Example: NaCl (~36g/100mL).
• Slightly soluble (sparingly soluble): Dissolves only to a limited extent; typically 0.1–1g per 100mL. Example: CaSO₄ (~0.2g/100mL), Ca(OH)₂ (~0.17g/100mL).
• Insoluble: Dissolves negligibly; typically <0.1g per 100mL. Example: BaSO₄ (~0.00024g/100mL), AgCl (~0.00019g/100mL).

Note: "Insoluble" doesn't mean zero solubility — even BaSO₄ dissolves a tiny amount!

This is due to lattice energy vs. hydration energy. Ba²⁺ is much larger than Mg²⁺, so BaSO₄ has a higher lattice energy (ions pack more tightly in the crystal). At the same time, the smaller Mg²⁺ has a much higher hydration energy (water molecules bind more strongly to it). For MgSO₄, hydration energy > lattice energy → soluble. For BaSO₄, lattice energy > hydration energy → insoluble. This trend continues down Group 2: BeSO₄ (very soluble) > MgSO₄ > CaSO₄ (slightly) > SrSO₄ > BaSO₄ (very insoluble).

Yes — nitrate (NO₃⁻) salts are universally soluble in water with no exceptions. This is because the nitrate ion is large, has a delocalized negative charge, and interacts weakly with cations, resulting in low lattice energies that are easily overcome by hydration. Even heavy metal nitrates like AgNO₃, Pb(NO₃)₂, and Ba(NO₃)₂ dissolve completely. This makes nitrate one of the most reliable ions in solubility predictions — if you see NO₃⁻, you can confidently mark the compound as soluble.

Only three cations form insoluble chlorides: Ag⁺ (silver), Pb²⁺ (lead(II)), and Hg₂²⁺ (mercury(I)). Their chlorides:

• AgCl — white precipitate, insoluble (Ksp ≈ 1.8×10⁻¹⁰)
• PbCl₂ — white precipitate, slightly soluble in cold water, more soluble in hot water
• Hg₂Cl₂ — white precipitate, insoluble (also called calomel)

All other common chlorides (NaCl, KCl, CaCl₂, MgCl₂, FeCl₃, AlCl₃, etc.) are freely soluble.

Temperature affects solubility differently for different compounds:

• Most solids: Solubility increases with temperature (e.g., KNO₃, NaCl, sugar).
• Some solids: Solubility changes very little (e.g., NaCl only increases from ~36g to ~39g per 100mL from 20°C to 100°C).
• Gases: Solubility decreases with increasing temperature (e.g., O₂, CO₂ in water).
• Special cases: PbCl₂ solubility increases dramatically with temperature — from ~0.9g at 20°C to ~3.3g per 100mL at 100°C.

The solubility rules on this page assume room temperature (~25°C) conditions unless otherwise noted.