Acid-base titration precision

In practice, however, any improvement in the sensitivity of an acid-base titration due to an increase in k is offset by a decrease in the precision of the equivalence point volume when the buret needs to be refilled. Consequently, standard analytical procedures for acid-base titrimetry are usually written to ensure that titrations require 60-100% of the buret s volume. 

The scale of operations, accuracy, precision, sensitivity, time, and cost of methods involving redox titrations are similar to those described earlier in the chapter for acid-base and complexometric titrimetric methods. As with acid-base titrations, redox titrations can be extended to the analysis of mixtures if there is a significant difference in the ease with which the analytes can be oxidized or reduced. Figure 9.40 shows an example of the titration curve for a mixture of Fe + and Sn +, using Ce + as the titrant. The titration of a mixture of analytes whose standard-state potentials or formal potentials differ by at least 200 mV will result in a separate equivalence point for each analyte. 

Acid-base titrations, especially at trace levels. Relative precision better than 1% at all levels. 

Common chemical titrations include acid-base, oxidation-reduction, precipitation, and complexometric analysis. The basic concepts underlying all titration are illustrated by classic acid-base titrations. A known amount of acid is placed in a flask and an indicator added. The indicator is a compound whose color depends on the pH of its environment. A solution of base of precisely known concentration (referred to as the titrant) is then added to the acid until all of the acid has just been reacted, causing the pH of the solution to increase and the color of the indicator to change. The volume of the base required to get to this point in the titration is known as the end point of the titration. The concentration of the acid present in the original solution can be calculated from the volume of base needed to reach the end point and the known concentration of the base. 

In a typical acid-base titration, the analyte is a base and the titrant is an acid, or vice versa. An indicator, a water-soluble dye (Section J), helps us detect the stoichiometric point, the stage at which the volume of titrant added is exactly that required by the stoichiometric relation between titrant and analyte. For example, if we titrate hydrochloric acid containing a few drops of the indicator phenolphthalein, the solution is initially colorless. After the stoichiometric point, when excess base is present, the analyte solution is basic and the indicator is pink. The indicator color change is sudden, so it is easy to detect the stoichiometric point. We then note the precise volume of titrant needed to cause the indicator to change color (Fig. L.3). 

The purpose of this chapter is to consider the preparation and standardization of acids and bases and to review some of the important applications of acid-base titrations in aqueous and nonaqueous systems. For end points to be detected most precisely, the pH in the vicinity of the equivalence point should change sharply. For this reason a solution of strong acid or base is chosen as titrant whenever possible. 

More than brief discussion of the numerous ways in which end points can be taken other than by visual methods is beyond our scope. For example, end-point techniques may involve photometry, potentiometry, amperometry, conductometry, and thermal methods. In principle, many physical properties can be used to follow the course of a titration in acid-base titrations, use of the pH meter is common. In terms of speed and cost, visual indicators are usually preferred to instrumental methods when they give adequate precision and accuracy for the purposes at hand. Selected instrumental methods may be used when a suitable indicator is not available, when higher accuracy under unfavorable equilibrium conditions is required, or for the routine analysis of large numbers of samples. 

As it was for acid-base titrations, the concept of relative precision (Section 3-7) is useful for comparing the steepness of titration curves in the immediate vicinity of the end point. For the formation of a precipitate of symmetrical charge type (m = n), with activity coefficients assumed to be unity. 

Finally we have seen in section 4-11 how acid-base titrations can be used in practice, even without any preliminary separations or sample clean-up, and what trade-offs are made in such analyses. This example illustrates a rather radical departure from the traditional emphasis on titrations as methods of high precision. As illustrated in Table 4.11-1, even when precise concentrations of well-defined chemical species cannot be derived from such complex mixtures, they nonetheless can be made to yield very useful quantitative information. 

Figure 5.15 shows a comparison of a low-frequency acid-base titration at two different ionic strengths with high-frequency titrations conducted at 3 and 10 MHz. In each case, 50 milliequivalents of HCl is titrated with 0.01 N NaOH. Obviously 10 MHz is the best frequency to use, but because of the curvature several additional titration points need to be taken to increase the precision of the endpoint determination. The M-shaped curve resulting at 3 MHz could lead to misinterpretation and an incorrect endpoint. 

More precisely, the color change of an indicator signals the end point of the titration, which if the proper indicator is chosen lies very near the equivalence point. Acid-base titrations are discussed in more detail in Section 17.3. 

As we found with acid-base titrations in the last chapter, however, the most precise method for locating the endpoint is the Gran method. This method results in a linear plot which intercepts the equivalence point at the X axis. Not only can we easily find the best line through linear regression but, as mentioned earlier, the necessary points can be taken at a distance from the equivalence point making this method rapid and convenient. 

An acid-base titration is a quick and convenient method for the quantitative analysis of substances with acidic or basic properties. Many inorganic and organic acids and bases can be titrated in aqueous media, but others, mainly organic, are insoluble in water. Fortunately, most of them are soluble in organic solvents hence they are conveniently determined by nonaqueous acid-base titrimetry. Although acid-base titrations can usually be followed potentio-metrically, visual endpoint detection is quicker and can be very precise and accurate if the appropriate indicator is chosen. 

For many years, analytical chemistry relied on chemical reactions to identify and determine the components present in a sample. These types of classical methods, often called wet chanical methods, usually required that a part of the sample be taken and dissolved in a suitable solvent if necessary and the desired reaction carried out. The most important analytical fields based on this approach were volumetric and gravimetric analyses. Acid-base titrations, oxidation-reduction titrations, and gravimetric determinations, such as the determination of silver by precipitation as silver chloride, are all examples of wet chemical analyses. These types of analyses require a high degree of skill and attention to detail on the part of the analyst if accurate and precise results are to be obtained. They are also time consuming, and the demands of today s high-throughput pharmaceutical development labs, forensic labs, commercial environmental labs, and industrial quality control... 

Neutralization indicators, or acid-base indicators or pH indicators, are auxiliary reagents added to the titrand solution in order to detect the equivalence point in acid-base titrations. They can also be used for an accurate quantitative measure of the pH. Tournesol, a natural pigment extracted from some blue-green lichens, was the first pH indicator to be used (1850). Phenolphthalein and methyl orange were introduced somewhat later (1877 and 1878, respectively). Undeniably, the chief interests in the use of acid-base indicators are their low cost and ease of handling. However, they give rise to less precise and less accurate endpoints than some instrumental methods. 

Precision of Acid-Base Titrations Related to the Sharpness Index... 

We cover acid-base titrations and indicators in more detail in Chapter 16. In most laboratory titrations, the concentration of one of the reactant solutions is unknown, and the concentration of the other is precisely known. By carefully measuring the volume of each solution required to reach the equivalence point, we can determine the concentration of the unknown solution, as demonstrated in Example 4.14. 

Another important application is the use of potentio-metric measurements for the evaluation of thermodynamic equilibrium constants. In particular, the dissociation constants for weak acids and weak bases in a variety of solvents are evaluated conveniently with a pH electrode measuring system. The most precise approach is to perform an acid-base titration such that the titration curve can be recorded. Obviously, one could measure the pH of a known concentration of a weak acid and obtain a value of its hydronium-ion activity which would permit a direct evaluation of its dissociation constant. However, this would be a one-point evaluation and subject to greater errors than by titrating the acid halfway to the equivalence point. The latter approach uses a well-buffered region where the pH measurement represents the average of a large number of data points. Similar arguments can be made for the evaluation of solubility products and stability constants of complex ions. 

Sodium hydroxide used to make standard NaOH(aq) solutions for acid-base titrations is invariably contaminated with some sodium carbonate, (a) Explain why, except in the most precise work, the presence of this sodium carbonate generally does not seriously affect the results obtained, for example, when NaOH(aq) is used to titrate HCl(aq). (b) Conversely, show that if Na2C03 comprises more than 1% to 2% of the solute in NaOH(aq), the titration results are affected. 

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