Weak bases curves

The relation between pH and the distribution ratio for a weak base (curve 3) is therefore the opposite of a weak acid so that it is possible to separate acids from bases in a mixture either by extracting the acids at low pH or the bases at high pH. Separating mixtures of acids or mixtures of bases is possible only if their dissociation or protonation constants differ by several pK units. 

It has been shown that for most acid-base titrations the inflection point, which corresponds to the greatest slope in the titration curve, very nearly coincides with the equivalence point. The inflection point actually precedes the equivalence point, with the error approaching 0.1% for weak acids or weak bases with dissociation constants smaller than 10 , or for very dilute solutions. Equivalence points determined in this fashion are indicated on the titration curves in figure 9.8. 

Potentiometric titration curves are used to determine the molecular weight and fQ or for weak acid or weak base analytes. The analysis is accomplished using a nonlinear least squares fit to the potentiometric curve. The appropriate master equation can be provided, or its derivation can be left as a challenge. 

Chemical Limitations to Beer s Law Chemical deviations from Beer s law can occur when the absorbing species is involved in an equilibrium reaction. Consider, as an example, an analysis for the weak acid, HA. To construct a Beer s law calibration curve, several standards containing known total concentrations of HA, Cmt, are prepared and the absorbance of each is measured at the same wavelength. Since HA is a weak acid, it exists in equilibrium with its conjugate weak base, A ... 

A slrang acid-weak base titration. The curve represents Ihe titration of 50.00 mL of 1.000 M ammonia, a weak base, with 1.000 M HCI. The solution at the equivalence point is acidic because of the NfV ion. Methyl red is a suitable indicator phenolphthalein would change color much too early. 

Weak acid and weak base. There is no sharp rise in the neutralisation curve and, generally, no simple indicator can be used. The titration should therefore be avoided, if possible. The approximate pH at the equivalence point can be computed from the equation ... 

Strong acid with a weak base. The titration of a strong acid with a moderately weak base (K sslO-5) may be illustrated by the neutralisation of dilute sulphuric acid by dilute ammonia solution [curves 1 and 3, Fig. 13.2(a)]. The first branch of the graph reflects the disappearance of the hydrogen ions during the neutralisation, but after the end point has been reached the graph becomes almost horizontal, since the excess aqueous ammonia is not appreciably ionised in the presence of ammonium sulphate. 

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 11.7 A typical pH curve for the titration of a weak base with a strong acid. The stoichiometric point (S) occurs at pH < 7 because the salt formed by the neutralization reaction has an acidic cation. 

Now consider the overall shape of the pH curve. The slow change in pH about halfway to the stoichiometric point indicates that the solution acts as a buffer in that region (see Fig. 11.3). At the halfwayr point of the titration, [HA] = [A ] and pH = pfCa. In fact, one way to prepare a buffer is to neutralize half the amount of weak acid present with strong base. The flatness of the curve near pH = pKa illustrates very clearly the ability of a buffer solution to stabilize the pH of the solution. Moreover, we can now see how to determine pKa plot the pH curve during a titration, identify the pH halfway to the stoichiometric point, and set pKa equal to that pH (Fig. 11.8). To obtain the pfCh of a weak base, we find pK3 in the same way but go on to use pKa -1- pfq, = pKw. The values recorded in Tables 10.1 and 10.2 were obtained in this way. 

We have already seen how to estimate the pH of the initial analyte when only weak acid or weak base is present (point A in Fig. 11.6, for instance), as well as the pH at the stoichiometric point (point S). Between these two points lie points corresponding to a mixed solution of some weak acid (or base) and some salt. We can therefore use the techniques described in Toolbox 11.2 and Example 11.6 to account for the shape of the curve. 

The principles that describe the titration of a weak acid also describe the titration of a weak base with hydronium ions. The titration curve for a weak base is shown in Figure 18-6. 

Schematic profile of the titration curve for a weak base B titrated with hydronium ions. The pH values are those for titration of ephedrine, a weak base that is the active ingredient in many decongestants.

Notice that the titration curve for a weak base has the same four regions seen in the titration curve of a weak acid ... 

C18-0146. The figure beiow shows the titration curves for 50.0-mL samples of weak bases A, B, and C. The titrant was 0.10 M HCl. (a) Which is the strongest base (b) Which base has the largest p (c) What are the initial concentrations of the three bases (d) What are the approximate p K- values of the conjugate acids of the three bases (e) Which of the bases can be titrated quantitatively using indicators Explain your answers, and identify an appropriate indicator for each base that can be titrated successfully. 

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. 

The absorption of short-chain weak acids in the rat intestine, as a function of pH, does not appear to conform to the pH partition hypothesis [44]. Similar anomalies were found with weak bases [77]. The apparent pKa values observed in the absorp-tion-pH curve were shifted to higher values for acids and to lower values for bases, compared with the true pKa values. Such deviations could be explained by the effect of an acid layer on the apical side of cells, the so-called acid pH microclimate [44,70,73,76-84],... 

For a weak acid, Pxw > Px and the log D curve decreases with pH for a weak base, Px > Pxh, and the log D curve increases with pH, according to this equation. 

Note that for a weak acid, the octanol causes the Bjerrum curve to shift in the direction of higher pH, whereas for a weak base, octanol causes the shift to lower values of pH. Equation (4.14) may be applied to the molecules in Fig. 4.7, and logP deduced from the shifts in the curves. 

Figure 6.8 shows the Bjerrum plots for an weak acid (benzoic acid, pKa 3.98, log So — 1.55, log mol/L [474]), a weak base (benzydamine, pKa 9.26, log So —3.83, log mol/L [472]), and an ampholyte (acyclovir, pKa 2.34 and 9.23, log So — 2.16, log mol/L I/40N ). These plots reveal the pKa and pA pp values as the pcH values at half-integral % positions. By simple inspection of the dashed curves in Fig. 6.8, the pKa values of the benzoic acid, benzydamine, and acyclovir are 4.0, 9.3, and (2.3, 9.2), respectively. The pA pp values depend on the concentrations used, as is evident in Fig. 6.8. It would not have been possible to deduce the constants by simple inspection of the titration curves (pH vs. volume of titrant, as in Fig. 6.7). The difference between pKa and pA pp can be used to determine log So, the intrinsic solubility, or log Ksp, the solubility product of the salt, as will be shown below.

It was found that the log 5/pH curves were altered in the presence of as little as 0.5% v/v DMSO, in that the apparent pKa values, pA pp, derived from log S versus pH [481], were different from the true pKa values by about one log unit. The pA pp values were generally higher than the true pKa values for weak acids (positive shift), and lower than those for weak bases (negative shift). This has been called... 

The Pm curve (dashed line) is not shifted to the right of the fraction neutral substance curve/ , (see inset in Fig. 7.34b). It just looks that way when unmatched scaling is used [554]. The two curves are exactly superimposed when the vertical coordinates of the Pm and / are normalized to a common value. The Pe curve, in contrast, is shifted to the right for weak acids and to the left for weak bases. In the log-log plot (log Pe vs. pH), the pH value at the intersection of the slope 0 and slope —1 curve segments indicates an apparent pKa (Fig. 7.34a). 

Figure 7.35 shows the characteristic log Pe-pH curve for a weak base, phenazo-pyridine (pKa 5.15). With bases, the maximum permeability is realized at high pH values. As in Fig. 7.34, the PAMPA assays were performed under iso-pH conditions (same pH in donor and acceptor wells), using the 2% DOPC in dodecane lipid... 

Membrane uptake of nonionized solute is favored over that of ionized solute by the membrane/water partition coefficient (Kp). If Kp = 1 for a nonionized solute, membrane permeability should mirror the solute ionization curve (i.e., membrane permeability should be half the maximum value when mucosal pH equals solute pKa). When the Kp is high, membrane uptake of nonionized solute shifts the ionization equilibrium in the mucosal microclimate to replace nonionized solute removed by the membrane. As a result, solute membrane permeability (absorption rate) versus pH curves are shifted toward the right for weak acids and toward the left for weak bases (Fig. 7). 

For weak acids, e.g., salicylic acid, the dependency on a pH gradient becomes complex since both the passive diffusion and the active transport process will be dependent on the proton concentration in the apical solution [61, 63, 98, 105] and a lowering of the pH from 7.4 to 6.5 will increase the apical to basolateral transport more than 20-fold. Similarly, for weak bases such as alfentanil or cimetidine, a lowering of the pH to 6.5 will decrease the passive transport towards the basolateral side [105]. The transport of the ionizable compound will, due to the pH partition hypothesis, follow the pKa curve. 

We can just sketch approximate titration curves of pH vs. percent of titration, since we do not have the concentration of the acid or the base, or the volume of solution being titrated. We can, however, precisely determine the pH at the half-equivalence point [mid-way between untitrated and completely titrated for a weak acid (or base)] this is equal to the pof the weak acid (or pOH = pATb of the weak base). Further, if we assume all solutions are 1.00 M, we can determine the pH at each equivalence point. For a weak base, the pH at the equivalence point equals -log >/0.50(ATa). For a weak acid, the pH at the... 

The effect of Kb on the titration curves for weak bases with a strong acid. 

The next point in the titration curve is the equivalence point. At this point, both the material added and the material originally present are limiting. At this point, neither of the reactants will be present and therefore will not affect the pH. If the titration involves a strong acid and a strong base, the pH at the equivalence point is 7. If the titration involves a weak base, only the conjugate acid is present to affect the pH. This will require a Ka calculation. If the titration involves a weak acid, only the conjugate base is present to affect the pH. This will require a Kb calculation. The calculation of the conjugate acid or base will be the moles produced divided by the total volume of the solution. 

The final region of the titration curve is after the equivalence point. In this region, the material originally present in the container is limiting. The excess reagent, the material added, will affect the pH. If this excess reactant is a weak acid or a weak base, this will be a buffer solution. 

It is a weak base, and the end point occurs at a pH between 4.5 and 5. It is a slightly stronger base than ammonium hydroxide, so its titration curve would appear similar to that of ammonium hydroxide in Figure 5.3(b). See Figure 5.10. 

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. 

If directed by your instructor, also obtain weak acid-strong base, strong base-strong acid, and weak base-strong acid titration curves. 

Figure 4.18 Conductimetric titration curves. As an acid is titrated with an alkali, so the ionic composition of the mixture changes and is reflected in the conductivity of the solution, (a) A strong acid and a strong base, (b) A strong acid and a weak base, (c) A weak acid and a weak base, (d) A weak acid and a strong base.

Abstract Titration of weak bases in non-aqueous solvents can provide valuable information about these weak bases. Some primary amines 1-aminobutane, 1-aminopropane, 2-aminoheptane, aminocyclohexane, 3-amino-l-phenylbutane were titrated with hydrochloric acid in toluene solvent. All the primary amines gave very well-shaped potentiometric titration curves. The same titrations were done with hydrochloric acid in methanol solvent to show the effect of amphiprotic solvent in the titrations with hydrochloric acid. 

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