Sodium hydroxide solution titration curve

When a saturated solution of sulphur dioxide is titrated against approximately 2 M sodium hydroxide solution the following pH curve is obtained Figure 10.4) ... 

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. 

Figure 4.6 A titration curve. Acetic acid (10 ml of a 0.1 mol l-1 solution) was titrated with a sodium hydroxide solution (0.2 mol l-1) and the pH of the resulting solution plotted against the amount of alkali added.

We apply the principle to compute the titration curve of 25 ml of 5M phosphoric acid with 0.1M sodium hydroxide solution. The pH-values in column A are given, the amount of the base solution to reach these pH-values are calculated in column J. 

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. 

Indirect Titration (Method B). A weighed sample of salt was added to a flask containing methanol and a known excess of aqueous sodium hydroxide solution. The mixture was stirred and warmed on a hot plate for 1 h. After cooling, the amount of excess hydroxide present was determined by titration with standard aqueous hydrochloric acid. The titration was monitored using a pH electrode and meter, and the end point was determined from the resulting titration curve. 

The complete titration curve is shown in the accompanying figure as a function of the amount X of sodium hydroxide solution added. Note that there is a very sharp pH change around the neutral (pH = 7) point. Consequently, the indicator used to determine the pH change does not have to be a very sensitive one. 

It is of interest to compare the titration curves for the weak acid with a strong base with the titration curve for a strong acid with a strong base. In Illustration 15.1-2 and here we have used equal amounts of acidic solutions of equal concentrations, and titrated both with the same sodium hydroxide solution. However, we see that the initial parts of the titration curve look somewhat different while the parts beyond neutrality are identical. In particular, the titration curve for the strong acid starts at a much lower pH than the titration curve for the weak acid. 

This convention is justified by its convenience, provided that (Section 1.4.2) there are no sudden inflection points in the neutralization curve of the must or wine at the pK of the organic acids present, as their buffer capacities overlap, at least partially. In addition to these somewhat theoretical considerations, there are also some more practical issues. An aqueous solution of sodium hydroxide is used to determine the titration curve of a must or wine, in order to measnre total acidity and buffer capacity. Sodium, rather than potassium, hydroxide is used as the sodium salts of tartaric acid are soluble, while potassium bitartrate would be likely to precipitate out during titration. It is, however, questionable to use the same aqueous sodium hydroxide solution, which is a dilute alcohol solution, for both must and wine. 

Soaps, sarcosinates and alkylether carboxylates (ethoxycarboxylates, polyethylene glycol monoesters) give poor curves in potentiometric titrations with benzethonium chloride, and ISO 2271 in alkaline solution does not always work very well for these compounds, particularly those of shorter chain length. The bromophenol blue method works well provided that the solution is distinctly alkaline it is necessary to add 5 ml 0.1 M sodium hydroxide before titrating. 

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... 

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. 

Titrations are often carried out in which one solution—either the analyte or the titrant—contains a weak acid or base and the other a strong base or acid. For example, we might be interested in the concentration of formic acid, a weak acid found in ant venom, and titrate it with sodium hydroxide. Alternatively, we might need to know the concentration of ammonia, a weak base, in a soil sample, and titrate it with hydrochloric acid. Figures 11.6 and 11.7 show the different pH curves that are found... 

Experimental measurements were carried out in several different solvents by two or three different methods. In some cases a shaking device was used and the rate of the reaction was followed by the evolution of carbon dioxide. In other cases the course of the reaction was followed by titrations using solutions which were sealed off in glass tubes immersed in thermostats and opened up under standardized sodium hydroxide or acid, as the case demanded. Many different tubes of the same solution were sealed off at once and each tube was used for a point on the time-concentration curve. The rate constant was determined from the slope of the line obtained by plotting the logarithm of the concentration against the time. 

Titrate the contents of each tube with 0.1 N sodium hydroxide, using bromothymol blue as the indicator. Disregard the results of the assay if the titration volume for the inoculated blank is more than 1.5 mL greater than that for the uninoculated blank. The titration volume for the 5.0-mL level of the Folic Acid Working Standard Solution should be approximately 8 to 12 mL. Prepare a standard curve by plotting the titration values, expressed in milliliters of 0.1IV sodium hydroxide for each level of the Folic Acid Working Standard Solution used, against the amount of folic acid contained in that tube. 

Three titrations of pure as well as NaCl- or Na2S04-containing nickel nitrate solutions with sodium hydroxide were carried out at 25°C. From the pH measured at each point taken on the titration curve (0.07 < (NaOH) / (Ni(N03)2) 1.62) the solubility product was calculated and extrapolated to zero ionic strength. A mean value of log,g K°f, = - 14.50 was obtained and has been taken as a basis for the value listed in Figure V-11 and Table V-6, respectively. 

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