Determination of end points in precipitation reactions

Many methods are utilised in determining end points in these reactions, but only the most important will be mentioned here. 

The theory of the process is as follows. This is a case of fractional precipitation (Section 2.8), the two sparingly soluble salts being silver chloride (Xsol 1.2 x 10 10) and silver chromate (Kso] 1.7 x 10 12). It is best studied by considering an actual example encountered in practice, viz. the titration of, say, 0.1M sodium chloride with 0.1M silver nitrate in the presence of a few millilitres of dilute potassium chromate solution. Silver chloride is the less soluble salt and the initial chloride concentration is high hence silver chloride will be precipitated. At the first point where red silver chromate is just precipitated both salts will be in equilibrium with the solution. Hence  

At the equivalence point [Cl ] = yjKso] AgC] = 1.1 x 10 5. If silver chromate is to precipitate at this chloride-ion concentration  

The difference is 1.3 x 10 5molL 1. If the volume of the solution at the equivalence point is 150mL, then this corresponds to 1.3 x 10 5 x 150 x 104/ 1000 = 0.02 mL of 0.1M silver nitrate. This is the theoretical titration error, and is therefore negligible. In actual practice another factor must be considered, viz. the small excess of silver nitrate solution which must be added before the eye 

The titration error will increase with increasing dilution of the solution being titrated and is quite appreciable (ca 0.4 per cent) in dilute, say 0.01 M, solutions when the chromate concentration is of the order 0.003-0.005M. This is most simply allowed for by means of an indicator blank determination, e.g. by measuring the volume of standard silver nitrate solution required to give a perceptible coloration when added to distilled water containing the same quantity of indicator as is employed in the titration. This volume is subtracted from the volume of standard solution used. 

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