Applications of Aqueous Acid-Base Titrations

Determination of acids For accurate results in the titration of weak acids with an indicator, one should be chosen that shows a transition color in the alkaline range coinciding as nearly as possible with the pH at the equivalence point. For best results a comparison solution of the indicator in a solution of the salt of the weak acid should be used. If the salt is not available, a buffer solution of the same pH may be substituted. 

For the titration of strong acids of concentration 0.1 M or higher with base containing some carbonate, the error involved in titrating to a pH of 4 is negligible. For the titration of more dilute solutions, however, it is advisable (Section 6-3) to titrate to the first perceptible color change of methyl red (pH range 4.4 to 6.0), boil to remove carbon dioxide, cool, and continue to the yellow color of the indicator. 

indicating that if an uncertainty of +0.1 pH unit is allowed a relative error of +1.2% may be expected. Even this precision can be attained only by use of a comparison standard of a pure bicarbonate solution containing the same concentration of indicator. Phenolphthalein, the most common indicator, does not become decolorized until a pH of 8.0, corresponding to an error of 3 to 5%. Color comparison is therefore essential. If thymol blue, a two-color indicator, is used, less attention need be paid to attaining identical indicator concentrations in the two solutions. By using a mixed indicator composed of thymol blue and cresol red, Simpson showed that results accurate to within 0.5% can be obtained without a comparison solution. 

FIGURE 6-2 Titration curves of mixed carbonates 0.15 to 1.S7 millimoles of Na2C03 and 0.05 to 0.21 millimoles of NaHCOs per liter. (From Cooper. ) 

Determination of hydroxyl ion in the presence of carbonate Owing to the small second ionization constant of carbonic acid K2 = 5 x 10 ), the titration of hydroxyl ion alone is not feasible in the presence of an appreciable concentration of carbonate ion. For example, a 0.01 M solution of sodium hydroxide has a pH of about 12, and a 0.01 M solution of sodium carbonate a pH of about 11.2. Thus the pH change in a titration of 0.01 M sodium hydroxide in the presence of 0.01 M sodium carbonate would be only about 0.8 pH unit. 

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