Titration error with acid/base indicators

Titration Errors with Acid/Base Indicators... 

It has been shown that for most acid-base titrations the inflection point, which corresponds to the greatest slope in the titration curve, very nearly coincides with the equivalence point. The inflection point actually precedes the equivalence point, with the error approaching 0.1% for weak acids or weak bases with dissociation constants smaller than 10 , or for very dilute solutions. Equivalence points determined in this fashion are indicated on the titration curves in figure 9.8. 

In Sections 10.11-10.16 it is shown how the change in pH during acid-base titrations may be calculated, and how the titration curves thus obtained can be used (a) to ascertain the most suitable indicator to be used in a given titration, and (b) to determine the titration error. Similar procedures may be carried out for oxidation-reduction titrations. Consider first a simple case which involves only change in ionic charge, and is theoretically independent of the hydrogen-ion concentration. A suitable example, for purposes of illustration, is the titration of 100 mL of 0.1M iron(II) with 0.1M cerium(IV) in the presence of dilute sulphuric acid ... 

We see that to obtain the most accurate results in titrating a strong acid and a strong base an indicator with indicator constant about 10 pK = 7) should be chosen, such as litmus or bromthymol blue. The titration curve calculated above, and given in Figure 20-3, shows however that the choice of an indicator is in this case not crucial any indicator with pK between 4 (methyl orange) and 10 (thymolphthalein) could be used with error less than 0.2%. 

When a solution of standard base is used only for titration of strong acids, a small amount of carbonate is not a serious source of error provided the end point is taken with an indicator that changes color at a pH of about 4 or 5. For standardizing such carbonate-containing solutions, potassium hydrogen phthalate is an unsuitable primary standard. An alternative is pure potassium chloride, which is passed through a cation-exchange column, converted to hydrochloric acid, and titrated with the sodium hydroxide. 

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. 

Figures 14-5 and 14-6 show that the choice of indicator is more limited for the titration of a weak acid than for the titration of a strong acid. For example. Figure 14-5 illustrates that bromocresol green is totally unsuited for titration of 0.1000 M acetic acid. Bromothymol blue does not work either because its full color change occurs over a range of titrant volume from about 47 mL to 50 mL of 0.1000 M base. An indicator exhibiting a color change in the basic region, such as phenolphthalein, however, should provide a sharp end point with a minimal titration error.

Titration the samples are titrated with a neutralizing agent and with a chemical indicator. The titration is conducted with an acid or a base which indicates the extent of the reaction. This analysis is simple, but subject to errors and inaccuracy and therefore should be repeated several times. 

It is suspected that an acid-base titrimetric method has a significant indicator error and thus tends to give results with a positive systematic error (i.e. positive bias). To test this an exactly 0.1 M solution of acid is used to titrate 25.00 ml of an exactly 0.1 M solution of alkali, with the following results (ml) ... 

The calculation of the titration error is directly connected to the preceding considerations. It is calculated as follows. Let fp be the potential at the final point (when the indicator color changes). The fp value is slightly different from the theoretical one ep (that at the equivalence point). Let (ppf be the fraction titrated at the final point. According to the definition of the titration error, already given when we considered acids and bases, the absolute titration error is %p — 1 since 1 is the fraction titrated at the equivalence point. The general expression of (p — 1, which is valid for any point of the titration curve, is (with i = 2)... 

The method described here is based on the difference between measurements of total alkaline earths by complexometric titration with EDTA (ethylenediamine-N,N,N, N -tetra-acetic acid) and selective measurement of calcium described in Section 11.2.1. The simultaneous EDTA titration of calcium, strontium and magnesium involves Eriochrome Black T (EBT) as indicator and was originally applied to seawater analysis by Voipio (1959) and Pate and Robinson (1961). To eliminate subjective errors in the determination of the endpoint, Culkin and Cox (1966) used photometric endpoint detection. A slight modification of this procedure, including the standardization of EDTA by magnesium is reported here. 

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