Titration curve of acetic acid

FIGURE 2-17 The titration curve of acetic acid. After addition of each increment of NaOH to the acetic acid solution, the pH of the mixture is measured. This value is plotted against the amount of NaOH expressed as a fraction of the total NaOH required to convert all the acetic acid to its deprotonated form, acetate. The points so obtained yield the titration curve. Shown in the boxes are the predominant ionic forms at the points designated. At the midpoint of the titration, the concentrations of the proton donor and proton acceptor are equal, and the pH is numerically equal to the pAfa. The shaded zone is the useful region of buffering power, generally between 10% and 90% titration of the weak acid. 

The titration curves of acetic acid, H2PO4, and NH4 (Fig. 2-18) have nearly identical shapes, suggesting that these curves reflect a fundamental law or relationship. This is indeed the case. The shape of the titration curve of any weak acid is described by the Henderson-... 

Figure 2-3. The titration curve of acetic acid. The molecular species that predominate at low and high pH are shown. At low pH (high [H ]), the molecule is proton-ated and has zero charge. As alkali is added, the [H+] decreases (H+ + OH" - H20), acetic acid dissociates, and the carboxyl group becomes negatively charged.

FIGURE 2-17 The titration curve of acetic acid. After addition of... 

Let us again take the example of acetic acid. Figure 1.7 represents the characteristic titration curve of acetic acid when it is titrated against a strong alkali. The figure traces the course of titration of a 0.1 N solution of acetic acid with 0.1 JV NaOH at 25 C. Before the titration is started (i.e. before any NaOH is added), the acetic acid is slightly ionized and the pH of the solution is due to acid alone. When successive aliquots of NaOH are added, the OH from dissociation of NaOH will combine with the free H in solution to form water. As soon as the free H is neutralized by OH" to water, some of the undissociated acetic acid immediately dissociates further to satisfy its dissociation constant. Thus with each addition of NaOH, more water is formed and more and more acetic acid gets converted to the acetate anion. 

FIGURE 2.12 The titration curve for acetic acid. Note that the titration curve is relatively flat at pH values near the pK in other words, the pH changes relatively little as OH is added in this region of the titration curve. 

Titration is the analytical method used to determine the amount of acid in a solution. A measured volume of the acid solution is titrated by slowly adding a solution of base, typically NaOH, of known concentration. As incremental amounts of NaOH are added, the pH of the solution is determined and a plot of the pH of the solution versus the amount of OH added yields a titration curve. The titration curve for acetic acid is shown in Figure 2.12. In considering the progress of this titration, keep in mind two important equilibria ... 

Thus the pK value is equal to the pH when the ratio of free acid to salt is unity, Lethe titration is half completed. The pK value for acetic acid in the titration represented in Fig. 2 is shown by the point x on the curve (b) and corresponds to a value of about pK =4.6 whereas the limiting thermodynamic value, pK, is 4.75. The estimate of PK may obviously be made at other points in the titration curve by computing the appropriate values of the ratios [A-]/[HA]. Or this, process may be reversed, and the curve (b) obtained from the observed value of pK and these ratios. The same procedure may also be extended to polybasic acids. A titration curve of malonic acid from the work of Gane and Ingold 5 is shown in Curve I of Fig. 3. Here... 

The titration curve for acetic acid shown in Figure 2-22 illustrates the effect of pH on the fraction of molecules in the un-ionized (HA) and ionized forms (A ). At one pH unit below the p/Q of an acid, 91 percent of the molecules are in the HA form at one pH unit above the pKg, 91 percent are... 

Figure 14 Conductometric titration curves of various acids by sodium hydroxide. Curve 1 represents a strong acid, and curve 5 an extremely weak one, while the others are intermediate. The acids are (1) hydrochloric acid, (2) dichloroacetic acid, (3) mon-ochloroacetic acid, (4) acetic acid, and (5) boric acid. (From Ewing GW (1985) Instrumental Methods of Chemical Analysis, 5th edn., p. 337. New York McGraw-Hill.)...

The shapes of the titration curves of weak electrolytes are identical, as Figure 2.13 reveals. Note, however, that the midpoints of the different curves vary in a way that characterizes the particular electrolytes. The pV, for acetic acid is 4.76, the pV, for imidazole is 6.99, and that for ammonium is 9.25. These pV, values are directly related to the dissociation constants of these substances, or, viewed the other way, to the relative affinities of the conjugate bases for protons. NH3 has a high affinity for protons compared to Ac NH4 is a poor acid compared to HAc. 

FIGURE 2.13 The titration curves of several weak electrolytes acetic acid, Imidazole, and ammonlnm. Note that the shape of these different curves Is Identical. Only their position along the pH scale Is displaced. In accordance with their respective affinities for ions, as reflected In their differing values. 

Weak acid with a strong base. In the titration of a weak acid with a strong base, the shape of the curve will depend upon the concentration and the dissociation constant Ka of the acid. Thus in the neutralisation of acetic acid (Ka— 1.8 x 10-5) with sodium hydroxide solution, the salt (sodium acetate) which is formed during the first part of the titration tends to repress the ionisation of the acetic acid still present so that its conductance decreases. The rising salt concentration will, however, tend to produce an increase in conductance. In consequence of these opposing influences the titration curves may have minima, the position of which will depend upon the concentration and upon the strength of the weak acid. As the titration proceeds, a somewhat indefinite break will occur at the end point, and the graph will become linear after all the acid has been neutralised. Some curves for acetic acid-sodium hydroxide titrations are shown in Fig. 13.2(h) clearly it is not possible to fix an accurate end point. 

Weak acids with weak bases. The titration of a weak acid and a weak base can be readily carried out, and frequently it is preferable to employ this procedure rather than use a strong base. Curve (c) in Fig. 13.2 is the titration curve of 0.003 M acetic acid with 0.0973 M aqueous ammonia solution. The neutralisation curve up to the equivalence point is similar to that obtained with sodium hydroxide solution, since both sodium and ammonium acetates are strong electrolytes after the equivalence point an excess of aqueous ammonia solution has little effect upon the conductance, as its dissociation is depressed by the ammonium salt present in the solution. The advantages over the use of strong alkali are that the end point is easier to detect, and in dilute solution the influence of carbon dioxide may be neglected. 

Beyond the stoichiometric point, in the final region of the titration curve, the concentration of acetic acid is very close to zero. There are no acid molecules to react with any further hydroxide ions, so excess hydroxide ions are... 

Attention is secondly focused on Figure 6.5 (B) which represents the titration curve of a weak acid against a strong base. The poor dissociation of the weak acid is reflected in the initial conductivity being low. The addition of alkali results in the formation of highly ionized sodium acetate and the conductance of the solution begins to increase. 

To select an indicator for an acid-base titration it is necessary to know the pH of the end point before using equation (5.5) or standard indicator tables. The end point pH may be calculated using equations (3.27), (3.29) or (3.30). Alternatively, an experimentally determined titration curve may be used (see next section). As an example, consider the titration of acetic acid (0.1 mol dm 3), a weak acid, with sodium hydroxide (0.1 mol dm-3), a strong base. At the end point, a solution of sodium acetate (0.05 mol dm 3) is obtained. Equation (3.28) then yields... 

FIGURE 5.2 A family of acid-base titration curves for a 0.10 M strong acid (HC1) and three weak acids, as indicated (0.10 M each), titrated with 0.10 M NaOH (strong base). HAc is a representation of acetic acid. 

In the process of a weak acid or weak base neutralization titration, a mixture of a conjugate acid-base pair exists in the reaction flask in the time period of the experiment leading up to the inflection point. For example, during the titration of acetic acid with sodium hydroxide, a mixture of acetic acid and acetate ion exists in the reaction flask prior to the inflection point. In that portion of the titration curve, the pH of the solution does not change appreciably, even upon the addition of more sodium hydroxide. Thus this solution is a buffer solution, as we defined it at the beginning of this section. 

Comparison of an alkalimetric titration curve of an equimolar (1 O 4 M) solution of acetic acid (pK = 4.8) and phenol (pK = 10) with a humic acid that contains 10 4 M carboxylic groups. 

In Investigation 8-A, you performed a titration and graphed the changes in the pH of acetic acid solution as sodium hydroxide solution was added. A graph of the pH of an acid (or base) against the volume of an added base (or acid) is called an acid-base titration curve. 

A plot of pH against the amount of NaOH added (a titration curve) reveals the pKa of the weak acid. Consider the titration of a 0.1 m solution of acetic acid (for simplicity denoted as HAc) with 0.1 m NaOH at 25 °C (Fig. 2-17). Two reversible equilibria are involved in the process ... 

Dissociation of the carboxyl group The titration curve of an amino acid can be analyzed in the same way as described for acetic acid. Consider alanine, for example, which contains both an a-carboxyl and an a-amino group. At a low (acidic) pH, both of these groups... 

Fig. 6.7. Conductometric titration curves of basic components extracted from engine oil CB SAE 30 into glacial acetic acid. Curve 1 -fresh oil, 2 - after 45 h of use, 3 - after 625 h of use (Pawlak et al.,...

Data in Table 3.2 can be represented in graph form, the titration curve. This is shown in Figure 3.1, and it is generated by plotting NaOH added versus pH. It shows a plateau approximately between pH 3.7 and 5.7 and an inflection point at pH 4.7, the pK of acetic acid. pH 3.7-5.7 is considered the acetic acid buffer buffering range. 

Curve (b) shows the titration of acetic acid, K, 1.75 x 10"1 mol dm"1 at two different concentrations. The initial additions of OH" establish a buffer solution in which the H30+ concentration is only slowly reduced. The resulting fall in conductance is increasingly counteracted by the addition of Nr and the formation of CHjCOO" thus leading to a minimum in the cur e. After the equivalence-point, there is a more rapid increase due to t -addition of excess OH" and Na+. Concentrated solutions of weak acids gi more pronounced changes of slope at the equivalence point than dilu solution. The change in slope is well defined for very weak acids, curve (c. 

Figure 6.19. Conductometric titration curves (a) Titration of strong acid with NaOH. (b) Titration of acetic acid, K, = 1.75 x 10 3 mol dm" with NaOH. (c) Titration of boric acid Jf,= 6x 10" io mol dm" J with NaOH. (d) Titration...

Figure 2-8. Titration curve of ethanoic (acetic) acid. AIM solution of ethanoic acid is titrated with NaOH. The Figure shows the dependence of the pH on the concentration of NaOH.

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