Strong base equivalence point

Theoretically, variations AB and ApH must be infinitely small, as the value of the AB/ ApH ratio at a fixed pH corresponds geometrically to the tangent on each point on the titration curve (Figure 1.4). More practically, buffer capacity can be defined as the number of strong base equivalents required to cause an increase in pH of 1 unit per liter of must or wine. It is even more practical to calculate smaller pH variations in much smaller samples (e.g. 30 ml). Figure 1.4 clearly shows the difference in buffer capacity of a model solution between pH 3 and 4, as well as between pH 4 and 5. 

Titrating a Weak Acid with a Strong Base For this example let s consider the titration of 50.0 mL of 0.100 M acetic acid, CH3COOH, with 0.100 M NaOH. Again, we start by calculating the volume of NaOH needed to reach the equivalence point thus... 

After the equivalence point NaOH is present in excess, and the pH is determined in the same manner as in the titration of a strong acid with a strong base. For example, after adding 60.0 mb of NaOH, the concentration of OH is... 

Figure 9.8b shows a titration curve for a mixture consisting of two weak acids HA and HB. Again, there are two equivalence points. In this case, however, the equivalence points do not require the same volume of titrant because the concentration of HA is greater than that for HB. Since HA is the stronger of the two weak acids, it reacts first thus, the pH before the first equivalence point is controlled by the HA/A buffer. Between the two equivalence points the pH reflects the titration of HB and is determined by the HB/B buffer. Finally, after the second equivalence point, the excess strong base titrant is responsible for the pH. 

Where Is the Equivalence Point We have already learned how to calculate the equivalence point for the titration of a strong acid with a strong base, and for the titration of a weak acid with a strong base. We also have learned to sketch a titration curve with a minimum of calculations. Can we also locate the equivalence point without performing any calculations The answer, as you may have guessed, is often yes ... 

The principal limitation to using a titration curve to locate the equivalence point is that an inflection point must be present. Sometimes, however, an inflection point may be missing or difficult to detect, figure 9.9, for example, demonstrates the influence of the acid dissociation constant, iQ, on the titration curve for a weak acid with a strong base titrant. The inflection point is visible, even if barely so, for acid dissociation constants larger than 10 , but is missing when is 10 k... 

Although not commonly used, thermometric titrations have one distinct advantage over methods based on the direct or indirect monitoring of plT. As discussed earlier, visual indicators and potentiometric titration curves are limited by the magnitude of the relevant equilibrium constants. For example, the titration of boric acid, ITaBOa, for which is 5.8 X 10 °, yields a poorly defined equivalence point (Figure 9.15a). The enthalpy of neutralization for boric acid with NaOlT, however, is only 23% less than that for a strong acid (-42.7 kj/mol... 

Perhaps the most obvious limitation imposed by Ks is the change in pH during a titration. To see why this is so, let s consider the titration of a 50 mb solution of 10 M strong acid with equimolar strong base. Before the equivalence point, the pH is determined by the untitrated strong acid, whereas after the equivalence point the concentration of excess strong base determines the pH. In an aqueous solution the concentration of H3O+ when the titration is 90% complete is... 

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... 

Thus, if we titrate a monoprotic weak acid with a strong base, the EW and FW are identical. If the weak acid is diprotic, however, and we titrate to its second equivalence point, the FW will be twice as large as the EW. 

The following data were collected with an automatic titrator during the titration of a monoprotic weak acid with a strong base. Prepare normal, first-derivative, second-derivative, and Gran plot titration curves for this data, and locate the equivalence point for each. 

A weak acid-strong base titration. The curve represents the titration of 50.00 mL of 1.000 M acetic acid, HC2H3O2. with 1.000 /W NaOH. The solution at the equivalence point is basic (pH = 9.22). Phenolphthalein is a suitable indicator. Methyl red would change color much too early, when only about 33 mL of NaOH had been added. Bromthymol blue would change color slightly too quickly. 

From Figure 14.7, it should be clear that the only suitable indicator listed is methyl red. The other two indicators would change color too early, before the equivalence point. In general, for the titration of a weak base with a strong acid, the indicator should change color on the acid side of pH 7. 

L C6H5NH3+ has 8.0 X 10-2 md particles (C6H5NH3+, Cl- ions) Or one could measure the pH at the equivalence point after titration with a strong base. 

It is clear that neither thymolphthalein nor phenolphthalein can be employed in the titration of 0.1 M aqueous ammonia. The equivalence point is at pH 5.3, and it is necessary to use an indicator with a pH range on the slightly acid side (3-6.5), such as methyl orange, methyl red, bromophenol blue, or bromocresol green. The last-named indicators may be utilised for the titration of all weak bases (Kb> 5 x 10-6) with strong acids. 

Weak acid and a strong base. The pH at the equivalence point is calculated from the equation ... 

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. 

In principle, any type of titration can be carried out conductometrically provided that during the titration a substantial change in conductance takes place before and/or after the equivalence point. This condition can be easily fulfilled in acid-base, precipitation and complex-formation titrations and also the corresponding displacement titrations, e.g., a salt of a weak acid reacting with a strong acid or a metal in a fairly stable complex reacting with an anion to yield a very stable complex. However, for redox titrations such a condition is rarely met. 

Considering the titration end-point or equivalence point, we in fact have to deal with the salt of a weak acid and a strong base or the reverse. Such a salt undergoes hydrolysis, e.g., for NaA, A + H20 HA + OH, so that the hydrolysis constant can be written as... 

In connection with the above, we shall still consider the pH curves of the displacement titrations, because in fact they represent a back-titration of an alkaline reacting salt (e.g., NaA) with a strong acid or of an acidic reacting salt (e.g., MX) with a strong base, so that in Fig. 2.19 the foregoing data of equivalence point and initial point are of direct application. 

Case C, the titration of a weak acid with a weak base and vice versa, has in fact already been illustrated in Fig. 2.18 by the curves BB and B B are fully valid and for characteristic (3) the initial point is still dependent on the original concentration c however for the further main part of the curve we see a clean symmetry versus the equivalence point, which has become a true inflection point, independent of the concentration and simply determined by the mean value of pKg and pKb, i.e., (p/ia + pKh)/2 or (pifa + pifw - pKa.)/2. It also means that in the simultaneous titration of a polyvalent acid or a series of weak acids of different strength with a strong base and vice versa, (1) the stronger the acid the earlier it is titrated within the series, (2) the initial point and the final end-point of the series are still influenced by the concentration, but (3) the intermediate steps are only determined at the pH of the inflection point by the mean value of the pifas of the subsequent acids and in its steepness by the difference between these pKgs. Therefore, consultation of pKa tables provides the most suitable way of predicting the results of such simultaneous titrations. 

Before the equivalence point, NaOH is the limiting reactant. The pH is determined from the concentration of excess HN03 remaining. The salt produced, NaN03, is the salt of a strong acid and a strong base. It will not affect the pH of the solution. 

After the equivalence point, the pH is determined directly from the concentration of excess NaOH since CH3COOH is now the limiting reactant. In the presence of the strong base, the effect of the weak base, CH3COO, derived from the salt is negligible. 

At the equivalence point, there is no excess acid or base. The pH of the salt solution, NH4NO3, is <7 since we have a solution of the salt of a weak base and a strong acid. The concentration of NH4NO3 determines the pH of the system. At point (g) ... 

At the equivalence point (pH = 7 at the equivalence point of a strong acid-strong base titration) ... 

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