Solubilities of ionic compounds depend in large measure upon the difference in the strength of the solute-solvent attractive forces and the solute-solute attractive forces. Because these differences are usually small, they do not vary in a regular way. Thus we use
solubility rules based on experimentally-determined solubilities.
A compound is classed as insoluble if the solubility is less than 0.1% by mass. Some compounds that are on the borderline of "soluble" and "insoluble" based on this criteria are indicated below.
Ionic compounds are soluble if the cation is ammonium (NH4+), sodium (Na+) and potassium (K+)
Ionic compounds are soluble if the anion is halide (Cl–, Br– and I–) except if the cation is Ag+ or Pb2+
nitrate (NO3–)
ethanoate (acetate CH3COO–)
perchlorate (ClO4–)
sulfate (SO42–) except BaSO4, PbSO4, SrSO4, CaSO4 (0.2%), Ag2SO4 (0.5%)
A general observation is that ionic compounds having anions with -1 charge tend to be soluble (with the exception of hydroxide as shown below).
Ionic compounds are INsoluble if the anion ishydroxide (OH–) or oxide (O2–) except if the cation is Na+, K+, Ba2+
carbonate (CO32–) except if the cation is Na+, K+, NH4+,
phosphate (PO43–) except if the cation is NH4+, Na+, K+
sulfide (S2–) except if the cation is Mg2+, Ca2+ , Sr2+, Na+, K+, NH4+.
A general observation is the ionic compounds that have anions with -2 or -3 charges are insoluble.