SOLUBILITY RULES
You should use the following rules to predict the solubility of ionic compounds in water. These rules
are only guidelines, and there are exceptions besides the ones listed. However, the ions normally
encountered in an introductory chemistry course, in most cases, will follow these rules.
This list of rules is in order of decreasing importance. That is, if the first rule is applicable then follow
it regardless of what the other rules may predict. Similarly, if the second rule is applicable then ignore
the third rule. Thus, you should only use the third rule when both rules one and two are not
applicable.
In general, these rules relate to the lattice energies of the particular compounds along with the
hydration energies of the individual ions. As a simplification, it is assumed that the greater the lattice
energy to lower the solubility. One of the factors determining the lattice energy are the magnitudes of
the charges on the ions (the greater the magnitude of the charges the greater the lattice energy). For
this reason, the basis of these rules is the magnitudes of the charges. These rules ignore other
influences on the lattice energies and the influence of the hydration energy. In addition, the rules
disregard the effect of temperature on solubility.
1. Compounds containing an ion with a +1 or -1 ion are normally soluble.
Examples: NaBr, K2SO4, FeCl2, AlCl3, Rb3PO4
Exceptions (Insoluble):
OH- (other than with cations that produce strong bases)
IB metals (column 11) and Hg22+
|(other than with NO3-|C2H3O2-, ClO3-, and
|ClO4-)
2+
2+
|
Pb , Hg with most -1 ions
2. Compounds containing an ion with a +3, -3 or higher charge ion are normally not soluble.
Examples: AlPO4, TiO2, Ca3(PO4)2
Exceptions (Soluble):
Cations combined with sulfate or dichromate.
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3. Compounds containing ions with a -2 charge are normally not soluble.
Examples: CaCO3, FeS, ZnC2O4
Exceptions (Soluble):
Dichromates
Sulfates (Insoluble only if combined with Pb, Hg, Ca, Sr, or Ba)
Note: The following ions are nearly always soluble:
Na+, K+, NH4+, NO3-, C2H3O2-, ClO3- and ClO4Note: Hydrides (H-) decompose in water to yield H2 and OHNote: Many metal oxides will react with water to produce hydroxides.
Examples:
NaCl
FeAsO4
Na3PO4
AgBr
Fe2(SO4)3
BaCO3
MgSO4
Ag2S
(NH4)2CO3
NaOH
AgOH
NaClO3
Hg2Cl2
Ba(OH)2
AgCl
NH4IO3
PbCl2
Ba(NO3)2
PbCrO4
ZnS
BaSO4
AgNO3
Ba(IO3)2
AgIO3
Soluble (Rule 1)
Insoluble (Rule 2)
Soluble (Rule 1)
Insoluble (Rule 1)
Soluble (Rule 2)
Insoluble (Rule 3)
Soluble (Rule 3)
Insoluble (Rules 1)
Soluble (Rule 1)
Soluble (Rule 1)
Insoluble (Rule 1)
Soluble (Rule 1)
Insoluble (Rule 1)
Soluble (Rule 1)
Insoluble (Rule 1)
Soluble (Rule 1)
Insoluble (Rule 1)
Soluble (Rule 1)
Insoluble (Rule 3)
Insoluble (Rule 3)
Insoluble (Rule 3)
Soluble (Rule 1)
Insoluble
Insoluble (Rule 1)
© 2006 Sevagram Enterprises
Na+ and ClFe3+ and AsO43Na1+
Ag+
SO42CO32SO42Ag+
NH41+
Strong base
Not a strong base
Na+ and ClO3Hg22+
Strong base
Ag+
NH4+ and IO3Pb2+
NO3CrO42S2Ba2+ + SO42NO3Exception
Ag+