Ionic strength chloride example

Most HPLC applications involving biomolecules utilize aqueous mobile phases. Critical parameters include both ionic strength and pH. Common solutes include TRIS, sodium phosphate, sodium acetate, and sodium chloride. Slightly alkaline pHs are preferable, for stability reasons. Specific examples of mobile phases include 50 mM TRIS, 25 mM KC1, and 5 mM MgCl2 (pH 7.2) for nucleotides, and 50 mM NaH2P04 (pH 7.0) and 20 mMTRIS and 0.1 M sodium acetate (pH 7.5) for both peptides and amino acids. All of these mobile phases are suitable for reverse phase or ion exchange applications. 

Such cases are not uncommon, but full quantitative treatments are rare, since often relatively large amounts of Y must be added to obtain measurable effects. Complications may then arise from the effects of the added Y on the nature of the medium (see Chapters 2 and 3). These are particularly notable when Y and I are charged, as is often the case. Under those circumstances, maintenance of the constant ionic strength of the medium with a known non-participating ionic species is essential. The classic case of common ion depression in solvolysis of benzhydryl chloride is dealt with in Chapter 2. A more recent example of this kind of treatment with neutral reactants occurs in the elucidation of the mechanism of olefin metathesis [20], catalysed by the ruthenium methylidene 9, Scheme 9.6. With ca. 5% of 9, disappearance of diene 10 was clearly not first order. However, reactions run in the presence of large excesses of phosphine 11 were much slower and showed first-order kinetics. The plot of kQ K against 1/ [ 11 ] was linear, consistent with dissociation of 9 to yield an active catalytic species prior to engagement with the diene, with k t [11] 3 > fc2[diene]. Because first-order kinetics were observed under these conditions, determination of order with respect to the catalytic species (as well as the diene) was simplified, and an outline for the mechanism could be constructed (see also Chapter 12 for more detailed consideration of catalysed olefin metathesis). 

Calibration techniques that use buffers to adjust the ionic strength and pH of the solution are effective for certain kinds of samples. For example, the detection of fluoride in public water supplies often is carried out by dilution of both standards and samples with a buffer that contains acetic acid, sodium chloride, and sodium citrate (with the pH adjusted to pH 5.0-5.5 by use of sodium hydroxide). This buffer performs three functions (1) it fixes the ionic strength of the standards and samples to the same level, principally determined by the buffer (2) the solution is buffered in a region where HO- ion does not interfere and (3) any Fe(III) or Al(III) ions are complexed by citrate to release the fluoride ion that is bound by these ions. 

By choosing a reasonable value of the ion size parameter a, independent of concentration, it is found that in many cases Eq. (3.126) gives a very good fit with experiment, often for ionic strengths up to 0.1. For example, on the basis of a = 0.4 nm, Eq. (3.126) gives an almost exact agreement up to 0.02 Min the case of sodium chloride (Fig. 3.34 and Table 3.10). 

A reduction in retention times can also be observed in analyses of solutions of high ionic strength. If gradient elution is initiated with a phosphate buffer with a concentration of, for example, 1 mM, and the content of sodium chloride is 0.1-... 

As indicated above, the e.m.f. of a cell with transference can be regarded as made up of the potential differences at the two electrodes and the liquid junction potential. It will be seen shortly (p. 229) that each of the former may be regarded as determined by the activity of the reversible ion in the solution contained in the particular electrode. In the cell depicted above, for example, the potential difference at the left-hand electrode is dependent on the activity of the chloride ions in the potassium chloride solution of concentration Ci similarly the potential difference at the right-hand electrode depends on the chloride ion activity in the solution of concentration Cz. For sufficiently dilute solutions the activity of a given ion, according to the simple Debyc-Huckel theory, is determined by the ionic strength of the solution and is independent of the nature of the other ions present. It follows, therefore, that the electrode potentials should be the same in all cells of the type... 

Other electrolytes in solution. For example, in a saturated solution of silver chloride in 0.01 M potassium chloride, the concentration of silver chloride is only about 10 M and so makes no appreciable contribution to the total ionic strength. Yet the activity coefficient of the silver ion is equal to that of potassium or chloride ions (0.89, from the DHLL). 

For cations and anions generally, the assumption that liquid-junction potentials are the same in the measurement of standards and unknowns is less likely to be valid than for pH measurements. It has been suggested that a quantity A ) expressed in pM or pA units be included in (13-26) and (13-27) to correct for changes injunction potential arising from differences in ionic strengths of standard and test solutions. Alternatively, these effects could be eliminated through the use of two reference half-cells composed of electrodes without liquid-junction potentials. For example, if the test solution contained chloride ion, both reference half-cells could be Ag/AgCl, and the liquid-junction potential would be eliminated. In practice, external reference half-cells without liquid junction are not always convenient. 

The solute solution concentration has been demonstrated to influence DRV encapsulation efficiency differently, depending on the solute. As an example, although glucose and CF entrapment values were found to decrease with increasing solute solution concentration, the same was not found true for encapsulation of sodium chloride and potassium chloride (1). For CF, best encapsulation yields in DRVs are demonstrated when a 17 mM solution in a tenfold dilution of an isotonic PBS buffer, is used. The ionic strength of the buffer used to dissolve the solute added at this step, should be at least 10 times less than that of the buffer used for DRV dilution after the hydration step (see below) in order to reduce material losses, due to osmotic activity of liposomes. 

When electrolyte concentrations are not too great, it is often useful to swamp both samples and standards with a measured excess of an inert electrolyte. The added effect of the electrolyte from the sample matrix becomes negligible under these circumstances, and the empirical calibration curve yields results in terms of concentration. This approach has been used, for example, in the potentiometric determination of fluoride ion in drinking water. Both samples and standards are diluted with a solution that contains sodium chloride, an acetate buffer, and a citrate buffer the diluent is sufficiently concentrated so that the samples and standaids have essentially identical ionic strengths. This method provides a rapid means of measuring fluoride concentrations in the part-per-million range with an accuracy of about 5% relative. 

The way to maintain the pH of the liquid solutions is to add an appropriate buffer system. This may also affect the overall stability of the formulation for example, the rate of deamidation appears faster in phosphate and bicarbonate buffers than in sulfate, nitrate, acetate, chloride, and pyruvate buffers (30). The ionic strength of the solution changes as the buffer concentration and other excipients, for example, salts to adjust the tonicity, are added, and this has an influence not just on the stability but also on the solubility (18). 

Optimize buffer conditions through testing buffers at different pH values at different concentrations and different ionic strengths, the latter being adjusted by the addition of sodium chloride (NaCI). Examples for a buffer optimization matrix are shown in Fig. 3. PTPs have peak activities usually between pH... 

Parenteral formulations often contain excipients considered to be chemically stable and inert however, all excipients in a formulation may influence the photochemical stability of the product. Dextrose and sodium chloride are used to adjust tonicity in the majority of parenteral formulations. Sodium chloride can affect photochemical processes by influencing solvation of the photoreactive molecules (see Section 14.2.3). The ionic strength is reported to affect the photochemical decomposition rate of minoxidil until a saturation level is reached (Chinnian and Asker, 1996). The photostability of L-ascorbic acid (vitamin C) in aqueous solution is enhanced in the presence of dextrose, probably caused by the scavenging effect of the excipient on hydroxyl radicals mediated by the photolysis of ascorbic acid sucrose, sorbitol, and mannitol have the same effect (Ho et al., 1994). Monosaccharides (dextrose, glucose, maltose, and lactose), disaccharides (sucrose and trehalose), and polyhydric alcohols (inositol, mannitol, and sorbitol) are examples of commonly used lyo-additives in parenterals. These excipients may also affect photochemical stability of the products after reconstitution. 

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