Poorly Buffered Solutions

Because they are weak acids or bases, the iadicators may affect the pH of the sample, especially ia the case of a poorly buffered solution. Variations in the ionic strength or solvent composition, or both, also can produce large uncertainties in pH measurements, presumably caused by changes in the equihbria of the indicator species. Specific chemical reactions also may occur between solutes in the sample and the indicator species to produce appreciable pH errors. Examples of such interferences include binding of the indicator forms by proteins and colloidal substances and direct reaction with sample components, eg, oxidising agents and heavy-metal ions. 

Slope too low Diaphragm contaminated/Clean diaphragm. Adsorption at glass membrane/Service glass membrane. Deswollen glass membrane after measurements in anhydrous solvents/Soak electrode in water between measurements. Old electrode/Regenerate glass membrane. Poor buffer solutions/Use fresh buffer solutions. 

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. 

The effects of solution acidity on the polarography of organic compounds have been reviewed, principally in aqueous solution. A thorough discussion of kinetic and catalytic currents that involve hydronium ion has been presented,52 and the irreversible polarographic and voltammetric curves that involve proton transfer in unbuffered and poorly buffered solutions have been discussed.59... 

It should be emphasized that even after indicator solutions have been neutralized they may still change the pH of poorly buffered solutions to a small extent. This point will be discussed in Chapter Ten. In most instances it is sufficient to prepare the... 

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. 

Solubility and stability of coelenterazine. Coelenterazine is very poorly soluble in neutral aqueous buffer solutions, and the solutions are unstable in air. It can be easily dissolved in water in the presence of alkali, but the resulting solution is extremely unstable under aerobic conditions. Coelenterazine is soluble in methanol, and the solution is relatively stable. The stability is enhanced by the addition of a trace of HCl. A methanolic solution of coelenterazine can be stored for several days at — 20°C, and a methanolic solution containing 1-2 mM HCl can be stored for several months at — 70°C under aerobic conditions without significant oxidation. In many other organic solvents, coelenterazine is less stable, and spontaneously auto-oxidized at significant rates. In dimethylformamide and DMSO, it is rapidly decomposed accompanied by the emission of chemiluminescence. e-Coelenterazines are generally less stable than coelenterazines. 

The poor solubility of coelenterazine in neutral aqueous buffer solutions often hampers the use of this compound in biological applications. The simplest way to make an aqueous solution is the dilution of a methanolic 3 mM coelenterazine with a large volume of a desired aqueous buffer solution. If the use of alcoholic solvents is not permitted, dissolve coelenterazine in a small amount of water with the help of a trace amount of 1 M NaOH or NH4OH, and then immediately dilute this solution with a desired aqueous buffer solution. However, because of the rapid oxidation of coelenterazine in alkaline solutions, it is recommended that the procedure be carried out under argon gas and as quickly as possible. 

A second form of storage iron is haemosiderin (Weir et al., 1984). This is deposited in humans as a response to the condition of iron overload. Haemosiderin forms as insoluble granules with electron dense cores surrounded by a protein shell. It exists in two forms primary haemosiderin is the result of iron overload due to excessive adsorption of iron in the gut, whereas the secondary form is caused by the numerous blood transfusions which are used to treat thallassaemia (a form of anaemia). Electron diffraction indicated that the iron core in primary haemosiderin is a 3-line ferrihydrite with magnetic hyperfine splitting only below 4 K and, in the secondary form, consists of poorly ordered goethite. As goethite is less soluble in ammonium oxalate buffer solution (pH 3) it has a lower intrinsic toxicity (Mann et al., 1988). 

An aqueous medium, either water or a buffered solution preferably not exceeding pH 6.8, is recommended as the initial medium for development of an IVIVC. Sufficient data should be submitted to justify pH greater than 6.8. For poorly soluble drugs, addition of surfactant (e.g., 1% sodium lauryl sulfate) may be appropriate. In general, nonaqueous and hydroalcoholic systems are discouraged unless all attempts with aqueous media are unsuccessful. Appropriate review staff in CDER should be consulted before using any other media. 

For a typical value of [NH4+]—say, 1.0 M—the NH3 concentration would have to be very small (0.0056 M). Such a solution is a poor buffer because it has little capacity to absorb added acid. Also, because the [NH3]/[NH4+] ratio is far from unity, addition of a small amount of H30+ or OH- will result in a large change in pH. 

Reconstituted acid-soluble collagen from various mineralized and unmineralized tissues have been examined for their potential to pick up calcium and phosphate from buffered solutions, and for their capacity to induce nucleation of a mineral phase426-434. Some collagens were good, others poor catalysts428,429 and apatite deposition proceeded in the presence of soft as well as of hard tissue collagens. The uptake of calcium ions requires the presence of phosphate ions and vice versa the Ca/P ratio is close to that of apatite (1.5—1.8)431. Study of exchange reactions by isotope tracers between solvent und substrate revealed that in absence of either... 

In all circumstances, as mentioned earlier, testing the solubility of the aqueous buffer solution in the mobile phase is crucial. Certain buffering salts have very poor solubility with solutions above 80% acetonitrile. When in doubt, test buffer solubility off-line using a small volume in vials and beakers on the bench top. Similarly, the sample should be soluble in the mobile phase. This is especially important if the sample was dissolved in a solvent which is different from the mobile phase. In this case, the only answer is to test the solution by injecting approximately 50 /u,L of sample solution into a vial containing approximately 1 mL of mobile phase and observing if the sample stays in solution. 

Many natural aqueous systems have ion combinations in solution that allow them to function as buffers. Other aquatic systems (e.g., some rivers and lakes) lack such ions or have a poor buffering capacity, and therefore they are vulnerable to acid or basic inputs. 

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