Electrometric determination

Polyaniline-modified electrodes allow electrometric determination of hydrogen peroxide produced in aminooxidase systems, without interference of electroreactive amino acids, such as cysteine, histidine, methionine, tyrosine and tryptophan359. 

Some of the volumetric methods described above may also be adapted to the electrometric determination of arsenic. Such methods have been described for titration of arsenites with ceric sulphate,9 iodine in the presence of sodium bicarbonate,10 chloramine (p-toluene-sulphone chloramide),11 alkaline potassium ferricyanide solution,12 potassium bromate13 or potassium iodate14 in the presence of hydrochloric acid, silver nitrate15 (by applying a secondary titration with 01N alkali to maintain the desired low H+-ion concentration), and with... 

The fact that the ionic product of D20 is several times smaller than the ionic product of ordinary water was one of the earliest firm results in deuterium chemistry (Abel et al., 1935 Wynne-Jones, 1936 Schwarzenbach et al., 1936). Three recent electrometric determinations (Gold and Lowe, 1963,1967 Covington et al., 1966 Pentz and Thornton, 1967) agree that the difference in pAw values is in the range 0-860-0-870 at 25°C a lower difference (0-809) was reported by Salomaa et at. (1964). 

Any phase boundary which responds to an ion concentration in a reproducible manner according to the Nernst equation can be used for the electrometric determination of that ion (S18). A wide variety have been described (B3, P13, R14, S18), but only those which have been successfully applied to biological systems will be considered here. 

Chikui, S. Electrometric determination in chromatographic development. III. Determination of alkaline earth metals during ascending development on paper, and their separation from alkali metal ions. Jap, Analyst 20, 167 (1971) Anal. Abstr. 23, 93 (1972)... 

Electrometrical determination of the redox potential (redox voltage)... 

If disturbing substances are absent, this process is also suitable as a calibrating method, e.g. as a check on the electrometric determination of oxygen. 

Earlier on we have seen that indicators are unsatisfactory for measuring pH of solutions which are colored. We have also seen that for works which require absolute accuracy, indicators are not a good tool for pH measurement. Such situations call for electrometric determination of pH. For a student of biochemistry, probably the most familiar instrument is the pH-meter. Its operation is quite simple it consists of a glass electrode which when dipped into a solution develops an electrical potential depending upon the hydrogen ion concentration and this potential is read off the display which is calibrated into a pH scale. Let us try to understand the basic principles underijdng the working of a pH meter. 

If we can fix the potential of the second electrode then the potential of other half-cells can be determined relative to it. In other words, the problem of measuring the potential of a half-cell in the presence of another half-ceU is solved by referring all potentials to a standard reference potential. The electrode(s) which is used as a standard is known as the reference electrode. We will consider some of the reference electrodes used in electrometric determination of pH in the following pages. 

Still the electrometrically determined intracellular [Cl ] of 18.7 mM is too high and too far from an equilibrium distribution to be consistent with passive transport in the three-compartment system. Thus intracellular chloride is at a higher electrochemical potential than either luminal or peritubular fluid chloride. This results in the passive efflux of intracellular Cl across the luminal and peritubular cell membranes. Since the tubule cell has greater [Cl ] than that expected from a simple passive distribution (5-7 mM) between intracellular and the two extracellular fluid compartments. Cl must be actively transported into the cell. 

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