E2 DISSOLUTION AND PRECIPITATION
Solubility rules
Ionic Compounds
When elements combine by transferring electrons from atoms of one element to those of another, an ionic compound is
formed; for example, NaCl (Na+ ions and Cl- ions) and BaI2 (Ba2+ ions and I- ions). An ionic compound, or salt, is
composed of oppositely charged ions packed into a crystal lattice and may include polyatomic ions in which a number
of atoms covalently bond together to form an ion; for example SrSO 4 (Sr2+ ions and SO42- ions) and NH4Cl (NH4+ ions
and Cl- ions).
Our description of ionic compounds represents an extreme and in practice there may be some covalent character
(electron sharing) in the bonds between the oppositely charged ions, for example Pb(CH3CO2)2 and HgCl2. When
writing the formula of a compound it is conventional not to indicate whether the compound is ionic or covalent. Thus,
the essentially ionic compound barium chloride is written BaCl2 and the essentially covalent compound water is written
H2O.
In addition, ionic compounds formed from aggregation of ions in a water solution may contain molecules of water
trapped in the crystal lattice, or in other words, water-of-crystallisation; for example CoSO4.H2O, CoSO4.6H2O and
CoSO4.7H2O. These exist in addition to the anhydrous compound, CoSO4. In most situations, any water-ofcrystallisation present is unimportant to the chemistry and any of the compounds may be used as a reagent: CoSO4,
CoSO4.H2O, CoSO4.6H2O or CoSO4.7H2O.
This is because cobalt(II) sulfate is water soluble and all compounds give rise to the same species in solution, Co2+(aq)
and SO42-(aq) where (aq) indicates the species is hydrated, ie surrounded by some water molecules.
The ions you will most commonly come across in the First Year Chemistry course are listed in Table E2. Note that
(i) ions are colourless in solution unless otherwise stated,
(ii) the table gives no indication of the relative stability of the ions given,
(iii) whereas the copper(I) ion is correctly represented as Cu+, the mercury(I) ion contains two mercury atoms
sharing a double charge and is correctly represented as Hg22+.
Solubility
Salts have widely differing water solubilities. At room temperature these range from very low values such as 10 -5 M for
BaSO4 and 10-17M for Ag2S to values around 10 M for some fluorides. It is convenient to use the following arbitrary
definitions:
soluble = solubility at least 0.1 M
slightly soluble = solubility between 0.01 and 0.1 M
insoluble = solubility less than 0.01 M
All salts are strong electrolytes and when the ions dissolve in water they do so with complete dissociation. This can be
represented by an equation, for example
KCl(s)  K+(aq) + Cl-(aq)
Na2SO4(s)  2Na+(aq) + SO42-(aq)
Note, that it would not be correct to represent these two solutions as KCl(aq) or as Na2SO4(aq), because any of the
compound that dissolves is present in solution as separated ions.
The process can be reversed if a solution containing the ions of a salt is evaporated to dryness, thus forming a solid; for
example
K+(aq) + Cl-(aq)  KCl(s)
The next step is to consider what happens when the solution of one salt is mixed with the solution of another. For
example silver nitrate dissolves in water, as does sodium chloride:
AgNO3(s)  Ag+(aq) + NO3-(aq)
NaCl(s)  Na+(aq) + Cl-(aq)
If the two solutions are mixed there will be a combination of silver ions, sodium ions, nitrate ions and chloride ions
present. Silver chloride is insoluble so as silver cations and chloride anions meet, they will combine to precipitate silver
chloride. The other combination of ions leads to sodium nitrate, which is soluble, and so these ions remain in solution
and are termed spectator ions. In chemistry, all reactions are represented by an ionic equation that indicates the change
that has occurred. Consequently, no spectator ions are included in the equation, as they have not undergone any change.
Ag+(aq) + Cl-(aq)  AgCl(s)
In this experiment, the solubility of a number of compounds is investigated and the application of these solubilities
studied. It is important to maintain clean equipment in order to prevent contamination and to use the recommended
quantities in order to get consistent results.
Hydroxides, oxides and carbonates
Although most hydroxides, oxides and carbonates are insoluble in water, they all react with acids as shown by the
following general equations. In these examples, M represents any metal ion, which for illustrative purposes has been
taken as divalent. The same type of reaction applies to the compounds of all metals.
M(OH)2(s) + 2H+(aq)  M2+(aq) + 2H2O(l)
MO(s) + 2H+(aq)  M2+(aq) + H2O(l)
MCO3(s) + 2H+(aq)  M2+(aq) + H2O(l) + CO2(g)
Let us look at these processes in a little more detail. Hydrochloric acid is a strong acid which means it is completely
dissociated in water into its constituent ions. It thus forms a solution that contains H+(aq) and Cl-(aq), and no HCl(aq).
It is the H+ that reacts with the carbonate to give water, carbon dioxide and release the metal ion into solution. The
chloride ion is merely a spectator ion; although it is present in the solution it takes no part in the reaction. If the water is
then evaporated the metal ion and the chloride ion combine to form a chloride salt. In this way one salt may be
converted into another by a series of chemical processes.
For example, to convert strontium carbonate to strontium iodide you might dissolve solid strontium carbonate in
hydroiodic acid and then evaporate the water from the resulting solution.
SrCO3(s) + 2H+(aq)  Sr2+(aq) + H2O(l) + CO2(g)
Sr2+(aq) + 2I-(aq)  SrI2(s)
Remember that the acids you use are aqueous solutions so if a salt dissolves in water, it would be expected to dissolve
in the water of a dilute acid solution. In your observations, look for changes which occur - there are the obvious ones
such as a solid dissolving or bubbles of a gas being produced but also less obvious ones such as evolution of heat. These
observations can be used to infer that a chemical reaction has taken place.
Table E2 - Some Commonly Occurring Ions
CATIONS
ANIONS
Na+
sodium ion
Cl-
chloride ion
K+
potassium ion
Br-
bromide ion
Mg2+
magnesium ion
I-
Ca2+
calcium ion
OH-
hydroxide ion
Sr2+
strontium ion
CN-
cyanide ion
Ba2+
barium ion
NCS-
thiocyanate ion
Al3+
aluminium ion
NO2-
nitrite ion
Pb2+
lead(II) ion
NO3-
nitrate ion
Sn2+
tin(II) ion
SO32-
sulfite ion
Sn4+
tin(IV) ion
SO42-
sulfate ion
Cr3+
chromium(III) ion
blue-green
HSO3-
hydrogensulfite ion
Mn2+
manganese(II) ion
very pale pink
HSO4-
hydrogensulfate ion
Fe2+
iron(II) ion
pale green
CO32-
carbonate ion
Fe3+
iron(III) ion
variable
HCO3-
hydrogencarbonate ion
Co2+
cobalt(II) ion
pink
PO43-
phosphate ion
Ni2+
nickel(II) ion
green
HPO42-
hydrogenphosphate ion
Cu+
copper(I) ion
H2PO4-
dihydrogenphosphate ion
Cu2+
copper(II) ion
MnO4-
permanganate ion
Ag+
silver(I) ion
ClO4-
perchlorate ion
Zn2+
zinc ion
CrO42-
chromate ion
yellow
Cd2+
cadmium ion
Cr2O72-
dichromate ion
orange
Hg22+
mercury(I) ion
S2-
sulfide ion
Hg2+
mercury(II) ion
HS-
hydrogensulfide ion
NH4+
ammonium ion
CH3COO-
acetate ion
C2O42-
oxalate ion
S2O32-
thiosulfate ion
blue
iodide ion
purple
[Fe(CN)6]4-
hexacyanoferrate(II) ion
yellow
[Fe(CN)6]3-
hexacyanoferrate(III) ion
brown