Reverse titration

The complete results, up to the addition of 200 mL of alkali, are collected in Table 10.3 this also includes the figures for 0.1 M and 0.01 M solutions of acid and base respectively. The additions of alkali have been extended in all three cases to 200 mL it is evident that the range from 200 to 100 mL and beyond represents the reverse titration of 100 mL of alkali with the acid in the presence of the non-hydrolysed sodium chloride solution. The data in the table are presented graphically in Fig. 10.2. 

For the reverse titration (chloride into silver nitrate), tartrazine (four drops of a 0.2 per cent solution per 100 mL) is a good indicator. At the end point, the almost colourless liquid assumes a blue colour. 

For the reverse titration (bromide into silver nitrate), rhodamine 6G (10 drops of a 0-05 per cent aqueous solution) is an excellent indicator. The solution is best adjusted to 0.05M with respect to silver ion. The precipitate acquires a violet colour at the end point. 

For reverse titration, i.e., of arsenite with hexacyanoferrate(III), the situation will be the same as in Fig. 3.79 the part of the curve beyond the end-point is a hyperbola again, for which Kies gave the following equation ... 

The reverse titration of thiosulphate with iodine is depicted in Fig. 3.81 and is often called the "reversed dead-stop end-point method . 

In fact, this has already been illustrated in Fig. 3.73 for the differential electrolytic potentiometric titration of Ce(IV) with Fe(II), both being reversible systems. This technique can be usefully applied, for instance, to the aforementioned KF titration of water and its reverse titration (cf., Verhoef and co-workers preference for bipotentiometric detection) in these instances the potentiometric dead-stop end-point titration and the reversed potentiometric dead-stop end-point titration, respectively, yield curves as depicted in Fig. 3.83. 

Bacterial cell walls contain different types of negatively charged (proton-active) functional groups, such as carboxyl, hydroxyl and phosphoryl that can adsorb metal cations, and retain them by mineral nucleation. Reversed titration studies on live, inactive Shewanella putrefaciens indicate that the pH-buffering properties of these bacteria arise from the equilibrium ionization of three discrete populations of carboxyl (pKa = 5.16 0.04), phosphoryl (oKa = 7.22 0.15), and amine (/ Ka = 10.04 0.67) groups (Haas et al. 2001). These functional groups control the sorption and binding of toxic metals on bacterial cell surfaces. 

A similar strategy has been applied by Mazik and coworkers. The receptor is usually titrated with increasing amounts of the sugar, to evidence binding, to extract the values of the binding constants and to get the stoichiometry of the complexes formed. Reverse titrations and extraction experiments can also be performed. The structures of the receptors described are based on different scaffolds the above mentioned hexasubstituted tripodal benzene moiety and 3,3, 5,5 -tetrasubstituted... 

N(A—B)/W, where (A—B) represents the corrected wlume of ferric titrant, N is the actual normality of the latter, and W is the weight of the sample, in g, represented by the aliquot taken from the volumetric. The results are at least as accurate, and probably more reliable than those obtd by the nitrometer method described in the JAN-N-494 Spec mentioned in the previous section. The main advantage of this reverse titration method is chat the NGu has much less opportunity to decompose in the alkaline condition imposed by the buffer before being reduced by the Ti(III)... 

To mimic c T-[(NH3)2Pt(H20)2]2+ with the more rapidly exchanging Pd11 and to prevent isomerization, it is necessary to employ the complex of ethylenediamine (en), [(en)Pd(H20)2]2+, that through chelation is necessarily cis. Upon titration with standard base an endpoint is reached after the addition of only one equivalent of base at pH 7.5, but the reversible titration curve is flattened on the pH axis and cannot be fitted with the equilibrium expression for a simple deprotonation. It was proposed that the mono-hydroxo complex dimerizes to abinuclear dihydroxo-bridged dimer [9], The two reactions and their equilibrium constant expressions follow. 

Fig. 15.4. Titration data from Tuominen (1967) for Cladonia alpestris, depicted as a function of pH versus concentration of added titrant. The closed circles represent forward titration data, while open circles stand for reversed titration data points. The upper curve is a calculated titration curve in pure water. The shaded area denotes the extent of pH buffering capacity exhibited by the lichen, relative to a non-buffering solution of pure water.

Fig. 8. Hypothetical titration curves illustrating time-dependence and irreversibility. Curve 1 is an apparently time-independent curve, obtained by continuous titration, waiting several minutes for each successive pH reading. Curve 2 is the reverse titration curve, beginning at the acid end point. Curve 3 is the forward titration curve obtained by flow methods, each pH being measured on a freshly mixed solution within seconds of mixing. Curves 4 and 5 are obtained from freshly mixed solutions with longer time intervals between mixing and measurement. Curve 6 is the titration curve which one might speculatively draw to represent instantaneous titration of the native protein.

Protein Isoionic pH Range of pH for instantaneous and reversible titration" Reversible major conformational change within this range s... 

Fig. 13. Titration data for the three reversibly titrated phenolic groups of ribo-nuolease at 25°C and at three ionic strengths, plotted according to Eq. (14) (Tanford et al., 1955a).

FIGURE 15-2 Titration curves for the Fe(II)-KMn04 reaction Qett) and the Fe(II)-K2Cr207 reaction (right). The downward trending curve is for the reverse titration of dichromate with Fe(II). (From Smith. )... 

In the foregoing discussion the indicator has tacitly been assumed to come rapidly to equilibrium at each point of the titration curve. That this is an over-simplihcation is evident from a number of experimental observations. Kolthoflf and Sarver found that the oxidation of diphenylamine with dichromate is induced by the Fe(II)-dichromate reaction. The direct oxidation is so slow that the indicator blank is best determined by comparison of the visual with the potentiometric end point. With ferroin. Smith and Brandt and Stockdale foimd that the reverse titration, dichromate with iron, gave satisfactory results at sufficiently high acidities, whereas the direct titration failed because the indicator could not be oxidized. Here the oxidation seems to be slow and the reduction rapid because of the irreversible nature of the oxidant and the reversible nature of the reductant. 

Actually, a strongly acidic solution of iodine can be titrated if the thiosulfate is added slowly with vigorous stirring. In the reverse titration (thiosulfate with iodine), however, a weakly acidic solution must be used to avoid decomposition. A low concentration of acid tends to prevent appreciable air oxidation of iodide. 

Fig. 6. Titration curves of insulin, zinc-free (A) and with. 1 mole zinc per 11,500 g. B and C). C is the direct titration of zinc insulin, B the reverse titration with acid or base. A is completely reversible. Data of Tanford and Epstein (1954a, b).

Despite the possibility of the simultaneous realization of these processes, the potentiometric curves of acid-base titration of Mo03 have no distinctive features as compared with the conventional ones [145]. However, the reverse titration curve differs appreciably from the direct titration results. Using the latter data the equilibrium constant of the formation of dimolybdate in the molten KCl-NaCl eutectic was calculated in Ref. [145] as... 

The curve of the reverse titration contains an additional drop of e.m.f. (pO) that corresponds to a Mo03/Mo04 buffer system and allowed estimation of the equilibrium constant of the following reaction as K = 880 880. 

Reverse titration results in the formation of an adduct, whose composition was assumed by Combes and Tremillon [154] to be W30 o The equilibrium constant of the reaction... 

The curves from the direct and reverse potentiometric titration of metavanadate ions with the Lux bases NaOH and Na2C03 are presented in Fig. 1.2.14a. The neutralization process runs in two stages according to the equations (1.2.87) and (1.2.88), to which two small diffuse pO-drops (bends) in the potentiometric curves correspond. Nevertheless, these bends become apparent on the differential potentiometric titration curve, shown in Fig. 1.2.14b. The reverse titration results in two steps of neutralization, also, that confirms the reversibility of the process of the acid-base neutralization of vanadium(V) oxocompounds in molten Nal. Comparison with the corresponding dependencies for the molten equimolar KCl-NaCl... 

Delimarskii et al. reported investigations of the solubility of ZnO, MgO, NiO and SrO in the molten KCl-NaCl equimolar mixture at 700 °C [236], The solubility was studied by the method of potentiometric titration of the solutions containing 0.01 mol kg-1 of the corresponding chloride with KOH, and of KOH solutions with quantities of the studied chloride (the reverse titration). 

Upon titration of 10 mM enPd(OH2)2 with standard base an endpoint is reached at pH 7.5 after addition of one equiv base. We discovered, however, that the reversible titration curve flattens on the pH axis and cannot be fitted with an equilibrium expression for a simple deprotonation (11). We succeeded in fitting the titration curve to a combination deprotonation and dimerization process that yields a binuclear, dihydroxo bridged dimer. For this reaction the overall equilibrium constant was found to be = 10 3 M. 

This mechanism is consistent with a recent EPR study of Pt/Ti02 catalysts in which only one Ti3+ species (which is assumed to be close to the metal particle) of four different species characterized could be reversibly titrated by CO (76a). 

The magnitude of the break will depend on both the concentration of the acid and the concentration of the base. Titration curves at different concentrations are shown in Figure 8.2. The reverse titration gives the mirror image of these curves. The titration of 0.1 M NaOH with 0.1 M HCl is shown in Figure 8.3. The selection of the indicators as presented in the figure is discussed below. 

Prasad and Dey [1962PRA/DEY] performed potentiometric titrations of 0.005 to 0.05 M ThCl4 solutions with NaOH solutions of different concentrations from pH 3 to pH 11-12 and reverse titrations by adding 0.04 and 0.08 M ThCk to the NaOH test solutions at 27°C. Ionic strength was not kept constant and the calibration of the glass-calomel electrode system used to measure pH is not mentioned. The authors report only the qualitative results that precipitation occurs already at < 4 (which is concluded from the inflection points of the titration curves) and that precipitation is complete near pH 7, independent of the thorium concentration. No thermodynamic data can be derived from this study. 

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