Aim: To show that
similar equilibria can be obtained, whether the reactants or the products are initially
The experiment shown in
illustration E03 illustrates two further aspects of chemical equilibria : the
reversibility of the reaction whether starting from the reactants or the products and the
demonstrable presence of the reactants and the products in the equilibrium mixture (see
The overall reaction in water (indicated by aq) can be written as follows:
starting with the reactants
1 g of the ochre solid (indicated by s) iron(III) chloride is dissolved in 100 mL of
distilled water to give a pale ochre solution. 1 g of white potassium iodide crystals are
then added, whereupon the solution rapidly becomes dark brown, indicating that a chemical
reaction has taken place.
What has happened? In
discussing this question utilize illustration E04 as an overlay to show the dissociation
and redox reactions.
||Each reactant freely dissolves in water
whereupon it dissociates into its ions:
As iron(III) iodide and potassium chloride are both quite soluble in water,
no precipitation is to be expected.
However, it can be shown that iron(II) chloride and iodine are formed,
the iodine being responsible for the brown colour observed. Fe3+
ions oxidize the iodide ions in a redox-reaction as shown below:
writing only the so-called essential species, all dissolved in water:
Illustration E03 indicates that if iron(II) chloride and iodine are dissolved
in water, this mixture will attain an equilibrium mixture containing both
the reactants and the products, although their concentrations may well
be different resulting in possible colour differences.
Even were the mixed quantities to correspond stoichiometrically, the concentrations
and hence the colour of the solution might still be different.
Furthermore the solubility of iodine in water is known to be enhanced
in the presence of iodide ions due to the formation of the more soluble
tri-iodide-ions, which also give a brown colour in water.