Titration equilibration time

The pH-metric method, which also requires no phase separation, has been used to determine dmg-liposome partitioning [149,162,385-387], The method is the same as that described in Section 4.14, except that FAT-LUV-ET liposomes are used in place of octanol. SUV liposomes have also been used [385,386]. To allow for pH gradients to dissipate (Section 5.6) in the course of the titration, at least 5-10 min equilibration times are required between successive pH readings. 

The initial titration aliquots were added automatically on the basis of the rate of change of EMF, mode (i), and the resulting time between aliquot additions was usually shorter than the 10 second equilibration time allowed in the mode (ii) titrations described above. The time differences were especially significant when the transmittance change per 0.05 cm3 aliquot was small, for example when the hyamine was present in excess at high salt concentrations. This means that the mode (i) titrations are more influenced by kinetic effects and so the measured curves are less distinctive as may be seen by comparing the results at 1.46% salt in Figures 6 and 7. 

FIGURE 8 Titration of 35 cm brine and 10 cm chloroform against SDBS 230 seconds initial equilibration time and 10... 

An example of these effects is shown in Figure 9 where a mode (ii) surfactant titration has been performed in the absence of salt and allowing a 10 seconds equilibration time between each aliquot. This gives a more clearly defined equivalence point when compared to the mode (i) titration in Figure 5. 

Fig. 16.14 Electron titration curve of a soil with Sn(OH)2 as a reduc-tant. Equilibration time 14 days (Lindsay Sadiq, 1983 with permission).

Equilibration time. It takes time for an electrode to equilibrate with a solution. A well-buffered solution requires 30 s with adequate stirring. A poorly buffered solution (such as one near the equivalence point of a titration) needs many minutes. 

In order to compare the PMMA results with those obtained with the carboxylic acrylic latex, the concentration of surface carboxyls must be determined. Acid location analysis (5 was carried out for this purpose. Briefly, the latexes were titrated conductometrically with 0.1N NaOH followed by a titration of the aqueous phase from which the particles had been removed by centrifugation. The difference in the two titrations provided the distribution between surface and soluble acid. The deficit between the total acid thus determined and the concentration of acrylic acid used in the polymerization was termed "buried". Although some drift occurred in the conductance with time, an equilibration time of approximately 10 minutes per addition of sodium hydroxide was generally sufficient to yield stable readings. 

Many reactions are catalyzed by acid sites on the surface of the catalyst. Isomerization, polymerization, aromatiza-tion, and cracking are catalyzed by Lewis and/or Bronsted acid sites. The precise nature of these sites is open to debate however, intuitively one can use an alkaline material to titrate acid sites and hence determine the number of such sites present. Beses, such as n-butylamine, with a series of Hammett indicators have been used for titrating acid sites. However, the system must be free from water contamination and the catalyst must be colorless to enable one to note indicator color changes. Diffusion of the indicators into the porous network can be very slow and require long equilibration times. 

The actual in vitro measurements of thermodynamic solubility correspond to the idealized titration regime conditions only to some extent. The closest method seems to be a labor-intensive shake-flask experiment requiring relatively large amounts of a dry crystalline drag and long equilibration times. Moreover, each pH point requires a separate measurement in different buffer. Different buffers usually represent different ionic conditions and may lead to internal inconsistency of the solubility profile thus obtained. The buffer issue deserves an entire subsection of this review and will be discussed later. The real question is whether this investment of resources is worthwhile. The answer depends on who is asking the question. A drag-development... 

Babic et al. [2] collected some lEP and PZC of activated carbons reported in literature (five references, 12 samples). The PZC range from pH 2.2 to 10.4 and lEP range from pH 1.4 to 7.1. With a few materials for which both PZC and lEP were reported these two values did not match. These discrepancies are not surprising. Most studied materials were not purified from mineral constituents, thus the measurements were not carried out at pristine conditions. It should be emphasized that potentioraetric titrations of activated carbons result in continuous drift of the pH (even after equilibration times of many hours), which is due to slow diffusion in... 

Motekaitis and Martell Model Contrary to the Ohman model, Motekaitis and Martell (1984) found no need to invoke polynuclear Al-citrate species to describe their potentiometric titration data. Their minimalist model, consisting of Al(Hcit)+(aqr), Alcif ( 7). and Al(H icit) (aqr) (Table 10.4), was generated from titrations of Al-citrate solutions in 0.10 mol KCl, using cUts/AIts ratios of 1 1, 2 1, and 3 1 from pH 2 to the onset of Al precipitation (approximately pH 7.5). They employed 4-hour equilibration times and included in their computations the AP+(aqr) hydrolysis for all four monomers [A10H2 (a ) to Al(OH)4 (a )], as well as Al2(OH)2 +( q) and AlTOH)4°+(6 g) [with log values obtained from the compilations of Smith and Martell... 

KCl solutions using cit s/Al s ratios of 1 1 or 2 1 and equilibration times of < 1 h. Data analysis included the consideration of citric acid dissociation (Table 10.1) and the formation of monomeric AP+(a ) hydrolysis products (Baes and Mesmer, 1986). Two chemical models were employed to describe the titration data. One model included Al2(H icit)22 (a ) (Table 10.4), while the other included Al2cit3 (a ) at the expense of Al2(H icit)22 (ag). Although the authors considered the Al2(H icit)22 (zz ) model to provide a... 

Analytical determinations of metal speciation in natural waters involve additions of metals or ligands to water samples. Rapid coordination equilibrium is then assumed, and, certainly, measurements of metal speciation by metal titrations demonstrate that some complexation reactions are fast on an analytical time scale. In some cases the effect of equilibration time (Kramer, 1986 Coale and Bruland, 1988) or the possible error due to incomplete equilibration (Sunda and Hanson, 1987) have been examined. 

Fluorescence was measured with a Turner model 111 filter fluorometer. The excitation filter was a Corning 7-60 (365 nm primary wavelength). The emission filters were Wratten 65-A (495 nm primary wavelength) and 2-A (sharp-cutoff below 415 nm). A digital multimeter was connected to the recorder terminals of the fluorometer to provide digital readout. Fluorescence-quenching (FQ) titrations were performed in batches. Preliminary experiments indicated that quenching was independent of time (at least 26 hours) after 30 minutes. Equilibration times of 60 minutes were used. 

A difficulty in the use of glass pH electrodes, even in aqueous solutions, but especially in partly aqueous solutions, is that they are relatively slow to respond to changes in hydronium ion activity. Equilibration times of at least 1 min are generally needed for each data point in a titration that might have a total of more than 20 points. This has two consequences (a) potentiometric titrations are relatively time consuming and thus (b) workers can become too impatient to wait for the electrode response to fully stabilize, leading to further measurement errors. Automated measurement, such as with Sirius autotitration equipment, is able to measure about 25 pKa values per day. Faster alternatives to the pH electrode are needed to rapidly measure pKa values. 

Mfj> (THEORY) EPOXY EQUIV. WT. (THEORY) TITRATED EQUIV. WT. WT. PERCENT EQUILIBRATION CATALYST EQUILIBRATION TIME, (HOURS) AT 80°C... 

Traditional methods to measure ionization constants are based on potentiometric or spectrophotometric titration. These methods are inherently slow due to the equilibration time needed after each addition of titrant typically one titration takes 20 to 40 minutes, and as most samples are poorly soluble in water, one needs to perform three titrations in water-co-solvent mixtures and extrapolate back to aqueous. 

The titrimetry approach is subject to a number of possible interferences. For example, in aqueous samples that are enriched in particulate material (riverine, estuarine, sewage) filtration is recommended (typically through a 0.45 pm membrane filter) to remove any particulate surfaces that might contribute to the solution alkalinity. In addition, aqueous samples that contain high concentrations of surfactants (e.g., industrial, sewage) may need longer pH electrode equilibration times during the alkalinity titrations. 

Equilibration time. In a well-buffered solution with adequate stirring, equilibration of the glass with analyte solution takes seconds. In a poorly buffered solution near the equivalence point of a titration, it could take minutes. 

Instead of the continuous titrations, the relative surface charge can also be determined in batch experiments. This involves the preparation of a number of independent batches and measuring the pH in these or in the supernatants after a certain equilibration time. Compared to the continuous titrations, which are typically computer controlled there are most probably more sources of experimental errors in the batch titrations (simply because more handling is... 

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