Concentration on titration curves

The effect of concentration on titration curves and indicators for a strong acid (HC1) and a strong base (NaOH). 100 cm3 of HC1 is being titrated with NaOH of the same molarity in each case. 

In earlier chapters, we considered the effects of reactant concentrations and completeness of the reaction on titration curves. Here, we describe the effects of these variables on oxidation/reduction titration curves. 

Figure 6. Effect of ionic strength (numbers in the legend are concentrations of NaCl in mol/1 soluble is hypothetical curve without influence of electrostatic potential of surface, i.e. Z = 0) on titration curves of silica with a particle diameter of 15 nm, pKa =7, a number of dissociating groups in a particle n = ns/2. The curves were calculated from Eq. (9).

All results apply to 1 F N(CHj)jCl and 20° unless otherwise stated. The titration of methanolic hydrochloric acid with lithiummethylate shows the expected Nernst relation between potential of the hydrogen electrode and total concentrations of H in all pertinent forms and that of methylate, respectively pK = -log [Htl tOR ] 16,60 pK 16,05, 1 F LiCl. This permits the definition of pH-scales based on molar concentrations. The titration curves or proton acids like acetic acid (pK = 8,9 ), acetylacetone (pK a 11,8), ammoniumion (pK = 11,2), oxalic acid (pKj = 8,4, pK = 5,lj) and pyrocatechol (pK 15,4, pK = 13,2) were as expected and permitted to check the measuring techniques. 

The effectiveness of any amine for absorption of both acid gases is due primarily to its alkalinity. The magnitude of this factor is illustrated in Figure 2-5, which shows pH values on titration curves for approximately 2N solutions of several amines when they are neutralized with CO2. The curves were obtained by bubbling pure CO2 through the various solutions and periodically determining the concentration of the solution and pH. The curve for an equivalent KOH solution is included for comparison. The relatively smooth curves for the amines, as compared to the sharp breaks in the KOH curve, may be interpreted as an indication of the presence of non-ionized species during neutralization of the former compounds. 

When the concentrations of HA and A are equal, equation 9.9 reduces to = [HaO ]) ot pH = pKa. Thus, the piweak acid can be determined by measuring the pH for a solution in which half of the weak acid has been neutralized. On a titration curve, the point of half-neutralization is approximated by the volume of titrant that is half of that needed to reach the equivalence point. As shown in Figure 9.20, an estimate of the weak acid s piQ can be obtained directly from the titration curve. 

It was indicated that the original method can be extended on systems where two or three analytes can be determined from a single titration curve. The shifts DpH affected by j-th PT addition should be sufficiently high it depends on pH value, a kind and concentration of the buffer chosen and its properties. The criterion of choice of the related conditions of analysis has been proposed. A computer program (written in MATLAB and DELPHI languages), that enables the pH-static titration to be done automatically, has also been prepared. 

If a sample contains groups that can take up or lose a proton, (N//, COO//), then one must expect the pH and the concentration to affect the chemical shift when the experiment is carried out in an acidic or alkaline medium to facilitate dissolution. The pH may affect the chemical shift of more distant, nonpolar groups, as shown by the amino acid alanine (38) in neutral (betaine form 38a) or alkaline solution (anion 38b). The dependence of shift on pH follows the path of titration curves it is possible to read off the pK value of the equilibrium from the point of inflection... 

As the titration begins, mostly HAc is present, plus some H and Ac in amounts that can be calculated (see the Example on page 45). Addition of a solution of NaOH allows hydroxide ions to neutralize any H present. Note that reaction (2) as written is strongly favored its apparent equilibrium constant is greater than lO As H is neutralized, more HAc dissociates to H and Ac. As further NaOH is added, the pH gradually increases as Ac accumulates at the expense of diminishing HAc and the neutralization of H. At the point where half of the HAc has been neutralized, that is, where 0.5 equivalent of OH has been added, the concentrations of HAc and Ac are equal and pH = pV, for HAc. Thus, we have an experimental method for determining the pV, values of weak electrolytes. These p V, values lie at the midpoint of their respective titration curves. After all of the acid has been neutralized (that is, when one equivalent of base has been added), the pH rises exponentially. 

Clearly for titration purposes, it is low-dielectric constant conducting solutions which will be important, and addition of a suitable reagent to such a solution permits the plotting of a titration curve from which the end point can be deduced as described in Section 13.7. It should be noted that in view of the enhanced conductance in the high-frequency field, the maximum concentration of reagents is much smaller than with normal conductimetric titrations, and the maximum concentration will depend on the frequency chosen. It is found that... 

Fig. 16-4 pH sensitivity to SO4- and NH4. Model calculations of expected pH of cloud water or rainwater for cloud liquid water content of 0.5 g/m. 100 pptv SO2, 330 ppmv CO2, and NO3. The abscissa shows the assumed input of aerosol sulfate in fig/m and the ordinate shows the calculated equilibrium pH. Each line corresponds to the indicated amoimt of total NH3 + NH4 in imits of fig/m of cloudy air. Solid lines are at 278 K, dashed ones are at 298 K. The familiar shape of titration curves is evident, with a steep drop in pH as the anion concentration increases due to increased input of H2SO4. (From Charlson, R. J., C. H. Twohy and P. K. Quinn, Physical Influences of Altitude on the Chemical Properties of Clouds and of Water Deposited from the Atmosphere." NATO Advanced Research Workshop Acid Deposition Processes at High Elevation Sites, Sept. 1986. Edinburgh, Scotland.)... 

When plotted on a graph of pH vs. volume of NaOH solution, these six points reveal the gross features of the titration curve. Adding additional calculated points helps define the pH curve. On the curve shown here, the red points A-D were calculated using the buffer equation with base/acid ratios of 1/3 and 3/1. Point E was generated from excess hydroxide ion concentration, 2.00 mL beyond the second stoichiometric point. You should verify these additional five calculations. 

It must be realized that the acidity of an acidic solution, expressed by its pH, is a physico-chemical property, which in fact (see calculations on pp. 83-85) represents a resultant of the identity and concentration of the acid even the overall pH height of the titration curve is still influenced by the concentrations of a strong acid, but for a weak acid that curve height, especially its h.n.pH value, forms a fairly reliable identity indication. 

Thus the best approach for HTS purposes is to experimentally determine the effect of enzyme titration on the observed reaction velocity, and to then choose to run the assay at an enzyme concentration well within the linear portion of the curve (as in Figure 4.6). Again, the other details of the assay conditions can affect the enzyme titration curve, so this experiment must be performed under the exact assay conditions that are to be used for library screening. 

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.

A dead-stop titration curve is produced if Ag+ is titrated with a halide using a pair of identical silver electrodes. Only whilst both Ag+ and Ag are present will a current flow in the cell, and this is linearly related to the Ag+ concentration. Bi-amperometric titrations require only simple equipment but generally give poorer precision because the currents measured are not necessarily on the limiting current plateau. 

In titrating a suspension of a-FeOOH (6 g e- 120 m2 g-1 2 10 4 mol g-1 surface functional groups (=FeOHTOT)) in an inert electrolyte (10 1 M NaCICU) with NaOH or HCI (Cb and Ca= concentration of base and acid, respectively, added per liter), we can write for any point on the titration curve... 

The point of zero charge is the pH at which net adsorption of potential determining ions on the oxide is zero. It is also termed the point of zero net proton charge (pznpc). It is obtained by potentiometic titration of the oxide in an indifferent electrolyte and is taken as the pH at which the titration curves obtained at several different electrolyte concentrations intersect (Fig. 10.5). It is, therefore, sometimes also termed the common point of intersection (cpi). The pzc of hematite has been determined directly by measuring the repulsive force between the (001) crystal surface and the (hematite) tip of a scanning atom force microscope, as a function of pH the pzc of 8.5-8.S was close to that found by potentiometic titration (Jordan and Eggleston, 1998). This technique has the potential to permit measurement of the pzc of individual crystal faces, but the authors stress that the precision must be improved. 

---