Glass electrode test

A high liquid junction potential may arise from sample clogging, precipitation in the junction, or if the junction is allowed to dry. In order to determine if this is the source of difficulty, the junction resistance test described in Section 3.2.5.B may be implemented. 

Another possible malfunction of the reference electrode is a broken internal. This is most obvious when a calomel internal is broken, thus causing the filling solution surrounding it to turn gray. The reference potential in this case is quite different from that of other calomel electrodes. The reference potential of an internal 

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Two methods are used to measure pH electrometric and chemical indicator (1 7). The most common is electrometric and uses the commercial pH meter with a glass electrode. This procedure is based on the measurement of the difference between the pH of an unknown or test solution and that of a standard solution. The instmment measures the emf developed between the glass electrode and a reference electrode of constant potential. The difference in emf when the electrodes are removed from the standard solution and placed in the test solution is converted to a difference in pH. Electrodes based on metal—metal oxides, eg, antimony—antimony oxide (see Antimony AND ANTIMONY ALLOYS Antimony COMPOUNDS), have also found use as pH sensors (8), especially for industrial appHcations where superior mechanical stabiUty is needed (see Sensors). However, because of the presence of the metallic element, these electrodes suffer from interferences by oxidation—reduction systems in the test solution. 

So-called combination electrodes may be purchased in which the glass electrode and the saturated calomel reference electrode are combined into a single unit, thus giving a more robust piece of equipment, and the convenience of having to insert and support a single probe in the test solution instead of the two separate components. 

The Nernst equation shows that the glass electrode potential for a given pH value will be dependent upon the temperature of the solution. A pH meter, therefore, includes a biasing control so that the scale of the meter can be adjusted to correspond to the temperature of the solution under test. This may take the form of a manual control, calibrated in 0 C, and which is set to the temperature of the solution as determined with an ordinary mercury thermometer. In some instruments, arrangements are made for automatic temperature compensation by inserting a temperature probe (a resistance thermometer) into the solution, and the output from this is fed into the pH meter circuit. 

Logarithmic scale for expressing acidity or alkalinity of water (7.0 to 0 indicates increasing acidity 7.0 to 14 indicates increasing alkalinity). Measured by means of a glass electrode/reference electrode pair immersed in the water sample under test. The potential difference depends upon the pH which is then displayed on a pH meter (high input impedance, millivoltmeter). 

A glass electrode, a thin-walled glass bulb containing an electrolyte, is much easier to use than a hydrogen electrode and has a potential that varies linearly with the pH of the solution outside the glass bulb (Fig. 12.11). Often there is a calomel electrode built into the probe that makes contact with the test solution through a miniature salt bridge. A pH meter therefore usually has only one probe, which forms a complete electrochemical cell once it is dipped into a solution. The meter is calibrated with a buffer of known pH, and the measured cell emf is then automatically converted into the pH of the solution, which is displayed. 

The electrolytes Na", and Cl are second only to glucose in being the most frequently run hospital tests. Many clinical chemistry analyzers now contain an ISE module for electrolyte analysis. Most commonly the module will consist of a Na -glass electrode, a valinomycin/PVC electrode, a Ag/AgCl pellet or a quaternary ammonium ion/PVC electrode and a reference electrode. A selective electrode for the bicarbonate ion continues to elude workers in the field. An indirect measurement of HCOf must be made. The sample is usually reacted with acid to evolve carbon dioxide gas which is measured with a traditional Severinghaus type CO2 electrode. Alternatively, the sample is treated with base to convert HCO to CO3 and a carbonate ion-selective electrode is used In this manner, the complete primary electrolyte profile is obtained electrochemically. 

Alkalinity is measured by acid-base titration with methylorange or phe-nolphthalein as indicator. Phenolphthalein changes color at pH 8.3, whereas methylorange changes color at pH 4.3. At pH 8 the neutralization of the strong alkali ingredients like NaOH is essentially complete. Further reduction of the pH to 4 will also measure carbonates and bicarbonates. Colorimetric tests and glass electrode systems are used to determine pH. 

Most suitable would be the use of a perfectly NH4+ ion-selective glass electrode however, a disadvantage of this type of enzyme electrode is the time required for the establishment of equilibrium (several minutes) moreover, the normal Nernst response of 59 mV per decade (at 25° C) is practically never reached. Nevertheless, in biochemical investigations these electrodes offer special possibilities, especially because they can also be used in the reverse way as an enzyme-sensing electrode, i.e., by testing an enzyme with a substrate layer around the bulb of the glass electrode. 

The membrane of the glass electrode is blown on the end of a glass tube. This tube is filled with a solution with a constant pH (acetate buffer, hydrochloric acid) and a reference electrode is placed in this solution (silver chloride or calomel electrodes). During the measurement, this whole system is immersed with another reference electrode into the test solution. The membrane potential of the glass electrode, when the internal and analysed... 

The paH response is tested by means of the paH values as determined by the indicator method the electromodve force of the cell immersed in buffer solutions whose pan is known is measured and the pan is spectro-photometrically determined and then plotted against E (Fig. 11). It can be seen that for this ethylene glycol-glass electrode the practical response is in good agreement with the theoretical one between pan 2 and 9 and for -1-21, -Hi, and — 19°C. The reproducibility of the determinations, estimated by the use of two different assemblies of electrodes, is better than 1.0 mV and the uncertainty of the pon determination is estimated at 0.1 pan unit. 

L. C. Clark first suggested in 1956 that the test solution be separated from an amperometric oxygen sensor by a hydrophobic porous membrane, permeable only for gases (for a review of the Clark electrode see [88]). The first potentiometric sensor of this type was the Severinghaus CO2 electrode [150], with a glass electrode placed in a dilute solution of sodium hydrogenocarbonate as the internal sensor (see fig. 4.10). As an equilibrium pressure of CO2, corresponding to the CO2 concentration in the test solution, is established in the... 

Therefore, the ISE potential depends on the CO2 partial pressure with Nernstian slope. Contemporary microporous hydrophobic membranes permitted the construction of a number of gas probes, developed mainly by the Orion Research Company (for a survey see [143]. The most important among these sensors is the ammonia electrode, in which ammonia diffusing through the membrane affects the pH at a glass electrode. Other electrodes based on similar principles respond to SO2, HCN, H2S (with an internal S ISE), etc. The ammonia probe has a better detection limit than the ammonium ion ISE based on the non-actin ionophore. The response time of gas probes depends mostly on the rate of diffusion of the test gas through the microporous medium [77,143]. 

Two methods are used to measure pH elec trot net ric and ehetnical indicator. The most common is electrometric and uses the commercial pH meter with a glass electrode. This procedure is based on the measurement ol the difference between the pH of an unknown or test solution and that of a. standard solution. 

FIGURE 18.6 A glass electrode consists of a silver wire coated with silver chloride that dips into a reference solution of dilute hydrochloric acid. The hydrochloric acid is separated from the test solution of unknown pH by a thin glass membrane. When a glass electrode is immersed in the test solution, its electrical potential depends linearly on the difference in the pH of the solutions on the two sides of the membrane. 

The concentration of the acid itself is of little significance other than analytical, with the exception of strong acids in dilute aqueous solutions. The concentration of H+ itself is not satisfactory either, because it is solvated diversely and the ability of transferring a proton to another base depends on the nature of the medium. The real physical quantity describing the acidity of a medium is the activity of the proton au. The experimental determination of the activity of the proton requires the measurement of the potential of a hydrogen electrode or a glass electrode in equilibrium with the solution to be tested. The equation is of the following type [Eq. (1.7)], wherein Cis a constant. 

The pH response was initially tested at 90°C and one atmosphere pressure by immersing the sensor in standard pH buffer solutions. The EMF of the sensor was then compared with the EMF generated by a commercial Ross-type glass electrode. Figure 3 shows the EMF versus pH response of a representative tube, PSU-T1-18. The response over a pH range of 2.5 to 9.5 is linear and displays a slope which is 98% of the theoretical Nernstian slope for this temperature. 

Normally, glass electrodes must be soaked in water for a few minutes to a few hours for the electrode to develop a pH response. This allows hydration of the outer surface to take place with the formation of the hydrogen ion-selective sites. Similarly, in 90°C pH tests at Pennsylvania State University, it has been found necessary to allow the zirconia electrodes to soak for several hours before a pH response was observed. 

A more critical test of the same data is shown in Fig. 3 where the points now represent the difference between determined pisT-values and the predictions of equation (119). Generally very satisfactory agreement between the results of the two sets of potentiometric studies with the glass electrode and those of the two conductivity studies is seen to exist. There is virtually a random scatter of all these points about the prediction from equation (119). The kinetic data (Brescia and La Mer, 1938) also fit in reasonably well, although they show some larger deviations and less good reproducibility, especially for large values of n. 

The measurement of surface pH of paper has been applied extensively for quick determination of acidity or alkalinity or for screening purposes when a large number of tests must be made. The surface pH test methods may be required when the integrity of the paper must be retained. The surface pH can be measured by application of appropriate indicators or by use of a glass electrode of the required shape—i.e., a flat-head electrode. A number of detailed procedures have been described (6). 

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