Acid and base titrations

This chapter describes several Important applications of aqueous equilibria. We begin with a discussion of buffer chemistry, followed by a description of acid and base titration reactions. Then we change our focus to examine the solubility equilibria of inorganic salts. The chapter concludes with a discussion of the equilibria of complex Ions. 

Brensted-Lowry Acids and Bases and Conjugate Pairs Lewis Acids and Bases Titration and Neutralization Hydrolysis... 

Electroacoustic measurements are usually carried out in titration mode, and several papers report results of both acid and base titration. Substantial hysteresis (Figure 1.9) was found in [242,427,491-493]. Other studies in similar systems report negligible hysteresis [444,449,494,495]. Four cycles of titration in ESA measurements reported in [496] produced different values of the potential in the acidic range but a common IEP. Similar lEPs were observed in titrations of alumina from pH 2 to 12 and back, but in titrations from pH 1 to 12, the IEP was shifted to high pH [497]. Generally, more pronounced hysteresis is expected in titrations over a wider pH range. 

By the use of indicators - compounds which change colour as their environment changes from an acidic to a basic one (or vice versa) - the precise moment of neutralization can be determined. For example, a measured quantity of acid solution of unknown concentration is placed in a flask with an indicator (e.g. litmus). A solution of base of known concentration is run into the flask until, with the addition of one excess drop of base, the colour changes. The volume of base run in has been measured and it is now possible to calculate the concentration of the acid solution. The method of calculation can be found in any elementary chemistry book and depends on knowledge of the equivalent weights of acids and bases. Titrations are used in paint chemistry to determine the acid value of a resin (Chapter 12) and it will be seen from the definition of acid value that no knowledge of equivalent weights is required for this determination it is sufficient to know the concentration of the alkali solution. 

Figure 1>1a. Strong Acid and Base Titration Curve tor the Entire Operating flange...

Figure 1 1d. Weak Acid and Base Titration Curve with Flattened Areas (Plateaus)...

Titrimetric (volumetric) factors for acids and bases are given in Table 11.28. Suitable indicators for acid-base titrations may be found in Tables 8.23 and 8.24. 

This approach can be used to sketch titration curves for other acid-base titrations including those involving polyprotic weak acids and bases or mixtures of weak acids and bases (Figure 9.8). Figure 9.8a, for example, shows the titration curve when titrating a diprotic weak acid, H2A, with a strong base. Since the analyte is... 

Thus far we have assumed that the acid and base are in an aqueous solution. Indeed, water is the most common solvent in acid-base titrimetry. When considering the utility of a titration, however, the solvent s influence cannot be ignored. 

Equivalent Weights Acid-base titrations can be used to characterize the chemical and physical properties of matter. One simple example is the determination of the equivalent weighf of acids and bases. In this method, an accurately weighed sample of a pure acid or base is titrated to a well-defined equivalence point using a mono-protic strong acid or strong base. If we assume that the titration involves the transfer of n protons, then the moles of titrant needed to reach the equivalence point is given as... 

Thompson, R. Q. Identification of Weak Acids and Bases by Titration with Primary Standards, /. Chem. Educ. 1988, 65, 179-180. 

Upon strong chelation, aU. four protons are displaced and base titration resembles that of a typical strong acid at four times the equivalent concentration. This statement is in agreement with equation 19, which shows that pM can be large (low concentration of free metal) at low pH if iC is large (strong chelation). 

There is also evidence for stable 3,4-adducts from the X-ray analysis of 2-amino-4-ethoxy-3,4-dihydropteridinium bromide, the nucleophilic addition product of 2-aminopteridine hydrobromide and ethanol (69JCS(B)489). The pH values obtained by potentiometric titration of (16) with acid and back-titration with alkali produces a hysteresis loop, indicating an equilibrium between various molecular species such as the anhydrous neutral form and the predominantly hydrated cation. Table 1 illustrates more aspects of this anomaly. 2-Aminop-teridine, paradoxically, is a stronger base than any of its methyl derivatives each dimethyl derivative is a weaker base than either of its parent monomethyl derivatives. Thus the base strengths decrease in the order in which they are expected to increase, with only the 2-amino-4,6,7-trimethylpteridine out of order, being more basic than the 4,7-dimethyl derivative. 

Hydrogen was recognized as the essential element in acids by H. Davy after his work on the hydrohalic acids, and theories of acids and bases have played an important role ever since. The electrolytic dissociation theory of S. A. Arrhenius and W. Ostwald in the 1880s, the introduction of the pH scale for hydrogen-ion concentrations by S. P. L. Sprensen in 1909, the theory of acid-base titrations and indicators, and J. N. Brdnsted s fruitful concept of acids and conjugate bases as proton donors and acceptors (1923) are other land marks (see p. 48). The di.scovery of ortho- and para-hydrogen in 1924, closely followed by the discovery of heavy hydrogen (deuterium) and... 

As pointed out in Chapter 4, an acid-base indicator is useful in determining the equivalence point of an acid-base titration. This is the point at which reaction is complete equivalent quantities of acid and base have reacted. If the indicator is chosen properly, the point at which it changes color (its end point) coincides with the equivalence point To understand how and why an indicator changes color, we need to understand the equilibrium principle involved. 

From Figure 14.5 and Example 14.7, we conclude that any indicator that changes ccdor between pH 4 and 10 should be satisfactory for a strong acid-strong base titration. Bromthymol blue (BB end point pH = 7) would work very well, but so would methyl red (MR end point pH = 5) or phenolphthalein (PP end point pH = 9). 

From Figure 14.6 and Example 14.8, it should be clear that the indicator used in this titration must change color at about pH 9. Phenolphthalein (end point pH = 9) is satisfactory. Methyl red (end point pH = 5) is not suitable. If we used methyl red, we would stop the titration much too early, when reaction is only about 65% complete. This situation is typical of weak acid-strong base titrations. For such a titration, we choose an indicator that changes color above pH 7. 

The complete results, up to the addition of 200 mL of alkali, are collected in Table 10.3 this also includes the figures for 0.1 M and 0.01 M solutions of acid and base respectively. The additions of alkali have been extended in all three cases to 200 mL it is evident that the range from 200 to 100 mL and beyond represents the reverse titration of 100 mL of alkali with the acid in the presence of the non-hydrolysed sodium chloride solution. The data in the table are presented graphically in Fig. 10.2. 

The Bronsted-Lowry theory of acids and bases referred to in Section 10.7 can be applied equally well to reactions occurring during acid-base titrations in non-aqueous solvents. This is because their approach considers an acid as any substance which will tend to donate a proton, and a base as a substance which will accept a proton. Substances which give poor end points due to being weak acids or bases in aqueous solution will frequently give far more satisfactory end points when titrations are carried out in non-aqueous media. An additional advantage is that many substances which are insoluble in water are sufficiently soluble in organic solvents to permit their titration in these non-aqueous media. 

Diphenylcarbazide as adsorption indicator, 358 as colorimetric reagent, 687 Diphenylthiocarbazone see Dithizone Direct reading emission spectrometer 775 Dispensers (liquid) 84 Displacement titrations 278 borate ion with a strong acid, 278 carbonate ion with a strong acid, 278 choice of indicators for, 279, 280 Dissociation (ionisation) constant 23, 31 calculations involving, 34 D. of for a complex ion, (v) 602 for an indicator, (s) 718 of polyprotic acids, 33 values for acids and bases in water, (T) 832 true or thermodynamic, 23 Distribution coefficient 162, 195 and per cent extraction, 165 Distribution ratio 162 Dithiol 693, 695, 697 Dithizone 171, 178... 

SOLUTION Because each acid molecule provides two protons, if the concentrations of acid and base were the same, the titration would require a volume of base equal to twice the volume of acid. The volume of base added to reach the stoichiometric point is less than twice the volume of acid so we can expect that the acid is less concentrated than the base. Proceed as in Toolbox L.2. 

What Do We Need to Know Already This chapter develops the ideas in Chapters 9 and 10 and applies them to equilibria involving ions in aqueous solution. To prepare for the sections on titrations, review Section L. For the discussion of solubility equilibria, review Section I. The discussion of Lewis acids and bases in Section 11.13 is based on Section 10.2. 

STRONG ACID-WLAK BASE AND WEAK ACID-STRONG BASE TITRATIONS... 

Strong Acid-Weak Base and Weak Acid-Strong Base Titrations... 

Figures 11.6 and 11.7 show the different pH curves that arc found experimentally for these two types of titrations. Notice that the stoichiometric point does not occur at pH = 7. Moreover, although the pH changes reasonably sharply near the stoichiometric point, it does not change as abruptly as it does in a strong acid-strong base titration.

FIGURE 11.10 The stoichiometric point of an acid base titration may be detected by the color change of an indicator. Here we see the colors of solutions containing a few drops of phenolphthalein at (from left to right) pH of 7.0, 8.5, 9.4 (its end point), 9.8, and 12.0. At the end point, the concentrations of the conjugate acid and base forms of the indicator are equal... 

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