The solubility would then equal the concentration of either the Ag+ or Cl- ions. Therefore, solubility = [Ag+] = [Cl-] and since we know that the Ksp of Silver Chloride = 1.6 x 10-10 moles per liter. Thus [Cl-] = (1.6 x 10-10) ½ producing 1.3 x 10-10 moles/ liter.
The precipitate is allowed to form in a slightly acidic medium as this prevents the rapidly forming precipitation from containing any unwanted anions which form as co-precipitates with Silver in neutral pH environment.
Furthermore, due to the speed of the precipitation the Silver Chloride is precipitated as extremely small particles as they don't have time to form large crystals and this may cause them to pass through the filter paper later on. As a result, slow heating and stirring is induced to allow the precipitate to form more proper and larger crystals.
Precautions must be taken in order to reduce light exposure as much as possible to the precipitate; this is because Silver Chloride decomposes into Silver and Chlorine in the presence of light as shown by the equation below:.
AgCl(s) ï‚® Ag(s) + ½Cl2(g).
If the above reaction happens in air then the Chlorine gas will escape and cause the analytical results to be low. None the less if this photochemical decomposition happens in the presence of excess Silver ions in an aqueous solution the following additional reaction occurs:.
3Cl2(g) + 3H20(l) + 5Ag+(aq) ï‚® 5AgCl(s) + CIO3-(aq) + 6H+(aq).
If the above reaction happens then the analytical results will be too high as the Silver Chloride produced increases the mass of the precipitate causing the analytical results to be higher than the actual value.
Additionally, to minimize the presence of unwanted ions in the precipitate the mixture is rinsed with 100mL fresh water and three 5mL samples of Acetone, however, since some of the Silver Chloride salt maybe dissolved in water, it will be lost through the process of filtering and go right through the filter apparatus, the worst case of mass loss is shown in the calculations below using the Ksp value:.