Solubility of ionic solids in water

The solubility of ionic solids in water varies widely, and it is most common to identify which are soluble using a set of rules based on experimental observations. Rules that pertain to XmYn compounds (X = cation; Y = anion) are given below.

These rules are quite complicated, but it useful to make the general observation that compounds where the charge on the ions is higher tend to be insoluble. 

Why is this so?
One of the factors that determines whether a substance is soluble in a given solvent is the relative strengths of the possible interactions:
solute-solute (in pure solute) 
solvent-solvent (in pure solvent)
solute-solvent (in a solution)
 
For soluble substances the solute-solvent attractive forces are stronger than the attractive forces between solute particles or solvent particles.

For ionic solids where the ions are highly charged, the attractive forces between the ions in the lattice are stronger than for ions with lower charges; thus the difference between the solute-solute attractive forces and solvent-solute attractive forces are greater.

 
  1. All Group 1 compounds (containing Na+, K+, Li+ ....) are soluble.
  2. All ammonium (NH4+) compounds are soluble.
  3. All compounds containing nitrate (NO3–), hydrogen carbonate (HCO3–) and hydrogen sulfate (HSO4–) ions are soluble.
  4. Most compounds containing chloride (Cl–), bromide (Br–), or iodide (I–) ions are soluble.
    Exceptions: Halides of Ag+ and Pb2+
  5. Most hydroxides (OH–) are insoluble.
    Exceptions: Group 1 hydroxides (see 1 above) and barium hydroxide (Ba(OH)2).
    Calcium hydroxide, (Ca(OH)2), is slightly soluble.
  6. All carbonates (CO32–), phosphates (PO43–) and sulfides (S2–) are insoluble.
    Exceptions: Compounds containing Group 1 metal cations and/or ammonium ion. See 1 and 2 above.
  7. Most sulfates (SO42–) are soluble. 
    Exceptions:  Barium sulfate (BaSO4) and lead sulfate (PbSO4) are insoluble.
    Calcium sulfate (CaSO4) and silver sulfate (Ag2SO4) are slightly soluble.


 
Dr Sheila Woodgate
© The University of Auckland 2002