Solubility Chart — Ionic Compounds in Water
| Anion | Na⁺ / K⁺ / NH₄⁺ | Mg²⁺ | Ca²⁺ | Ba²⁺ | Fe²⁺ | Fe³⁺ | Cu²⁺ | Zn²⁺ | Ag⁺ | Pb²⁺ | Al³⁺ |
|---|---|---|---|---|---|---|---|---|---|---|---|
| NO₃⁻ (nitrate) | S | S | S | S | S | S | S | S | S | S | S |
| CH₃COO⁻ (acetate) | S | S | S | S | S | S | S | S | SS | S | S |
| Cl⁻ (chloride) | S | S | S | S | S | S | S | S | I | SS | S |
| Br⁻ (bromide) | S | S | S | S | S | S | S | S | I | SS | S |
| I⁻ (iodide) | S | S | S | S | S | S | S | S | I | I | S |
| SO₄²⁻ (sulfate) | S | S | SS | I | S | S | S | S | SS | I | S |
| OH⁻ (hydroxide) | S | I | SS | S | I | I | I | I | I | I | I |
| CO₃²⁻ (carbonate) | S | I | I | I | I | I | I | I | I | I | I |
| PO₄³⁻ (phosphate) | S | I | I | I | I | I | I | I | I | I | I |
| S²⁻ (sulfide) | S | S | S | S | I | I | I | I | I | I | S |
| CrO₄²⁻ (chromate) | S | S | S | I | S | S | S | S | I | I | S |
| C₂O₄²⁻ (oxalate) | S | SS | I | I | I | I | I | I | I | I | I |
Categories use the conventional general-chemistry thresholds: S = soluble (>0.1 mol/L at 25 °C), I = insoluble (<0.01 mol/L), SS = slightly soluble (between). These are operational labels, not Ksp values — for quantitative work, look up Ksp directly. Five rules cover most cases: (1) all NO₃⁻ salts soluble; (2) all Na⁺, K⁺, NH₄⁺ salts soluble; (3) Cl⁻, Br⁻, I⁻ soluble except with Ag⁺, Pb²⁺, Hg₂²⁺; (4) SO₄²⁻ soluble except BaSO₄, PbSO₄, SrSO₄, with CaSO₄ slightly; (5) OH⁻, CO₃²⁻, PO₄³⁻, S²⁻ insoluble except with Group 1 and NH₄⁺ (and Ba(OH)₂, BaS). Values at 25 °C. Source: CRC Handbook, Zumdahl's Chemistry.
Frequently Asked Questions
How do you use a solubility chart to predict precipitation?
Three steps: (1) write the four possible cation-anion swaps from the two reactants; (2) check each combination against the chart; (3) any product marked I (or SS, depending on concentration) precipitates. Example: NaCl(aq) + AgNO₃(aq). The two new combinations are NaNO₃ (soluble — every nitrate is) and AgCl (insoluble — chlorides fail with Ag⁺, Pb²⁺, Hg₂²⁺). White AgCl precipitates and Na⁺/NO₃⁻ remain spectator ions. The net ionic equation is Ag⁺(aq) + Cl⁻(aq) → AgCl(s).
Why are all nitrate salts soluble?
Nitrate is large (radius ~189 pm), polyatomic, and bears its single negative charge delocalized over three oxygens, so it forms only weak ion-ion attractions in the solid. The resulting low lattice energy is more than offset by favorable hydration energies of the cation and the nitrate itself, making dissolution exergonic for every cation tested. That universal solubility is why nitrate is the go-to counterion when you need an aqueous source of an unusual cation — Pb(NO₃)₂, AgNO₃, Hg(NO₃)₂ all dissolve cleanly.
What does 'slightly soluble' mean in practice?
SS compounds reach a limited equilibrium concentration before precipitating — somewhere between 0.01 and 0.1 mol/L. They give measurable ion concentrations and a visible solid only once that ceiling is exceeded. Examples: CaSO₄ (Kₛₚ = 4.9 × 10⁻⁵), PbCl₂ (Kₛₚ = 1.7 × 10⁻⁵), Ag₂SO₄ (Kₛₚ = 1.2 × 10⁻⁵). Whether a slightly soluble salt actually precipitates in a given mixture depends on Q vs. Kₛₚ. Temperature matters too: PbCl₂ dissolves about three times more in boiling water than at 25 °C, which separates it from AgCl in classical qualitative analysis.