Most ionic compounds dissolve in water, but a few groups follow clear patterns that make prediction easy.
General solubility rules
Salts containing alkali‑metal cations (Li+, Na+, K+, Rb+, Cs+) and the ammonium ion (NH4+) are virtually always soluble. Nitrates (NO3‑), acetates (CH3COO‑) and most perchlorates (ClO4‑) also dissolve without exception.
Halide exceptions
Chlorides, bromides and iodides are generally soluble, but silver (Ag+), lead(II) (Pb2+) and mercury(I) (Hg2 2+) form insoluble halides. Silver chloride (AgCl), silver bromide (AgBr) and silver iodide (AgI) precipitate readily. Lead(II) chloride (PbCl2) is only moderately soluble in cold water and becomes more soluble when heated. Mercury(I) chloride (Hg2Cl2) is also poorly soluble.
Sulfate exceptions
Most sulfates dissolve, yet barium sulfate (BaSO4), lead(II) sulfate (PbSO4) and strontium sulfate (SrSO4) are essentially insoluble. Calcium sulfate (CaSO4) is sparingly soluble, which explains its presence in hard water.
Carbonates, phosphates, sulfides and hydroxides
Salts of carbonate (CO3 2‑), phosphate (PO4 3‑), sulfide (S2‑) and hydroxide (OH‑) ions are usually insoluble, except when paired with alkali‑metal cations or ammonium. Transition‑metal compounds in these families (e.g., FeCO3, Cu3(PO4)2, ZnS, Ni(OH)2) are typically sparingly soluble.
Effect of pH on solubility
Acidic conditions increase the solubility of salts that contain basic anions. For instance, carbonate salts dissolve in acid because CO3 2‑ reacts with H+ to form CO2 and water, pulling the equilibrium toward dissolution. The same principle applies to phosphate and hydroxide salts.
Using solubility product constants (Ksp)
When a salt is only slightly soluble, its solubility can be quantified by the solubility product constant, Ksp. For silver chloride, Ksp = [Ag+][Cl‑]. A small Ksp indicates low solubility; a larger value means the salt dissolves more readily.