Glass electrode deviations

Figure 3.12 Demonstration of the extent of the alkaline error, showing plots of the deviations in pH measurement under aqueous alkaline conditions as determined with a standard glass electrode (the Coming 015 ). The figures in parentheses represent the concentrations of the metal hydroxide salt in mol dm". From Christian, G. D., Analytical Chemistry, 5th Edn, Wiley, 1994. Reprinted by permission of John Wiley Sons, Inc.

The thermodynamic stability constant of silver sulfadiazine in water has been measured by Boelema et al. (20). They made use of a microcomputer-controlled titrator and measured simultaneously the pH value with a Radiometer G2040C glass electrode and the silver ion concentration with an Orion research 941600 ion-selective silver sulfide electrode. The measured stability constant and standard deviation were logK=3.62i 0.05 n = 9, temperature 25°C, ionic strength 0.1 (sodium nitrate). The calculated conditional stability constant log K at pH 7.4 was 3.57. 

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. 

This conclusion was reinforced in another way. The observed emf for the glass electrode-calomel cell (with liquid junction) in eight solutions with pH 2-9 differed by a constant amount from that for the glass-AgCl Ag combination (without liquid junction) in the same cell. The range was from 9.0 mV in the acetate buffer to 9.6 mV in the tris buffer the mean was 9.43 mV, and the standard deviation was 0.19 mV. 

In practice, a glass electrode is almost always used in place of the Pt(H2) electrode. A glass electrode has a deviation from the H+ (aq) ion response function (non-Nernstian behavior) and, therefore, should be calibrated using a set of the standard (buffer) solutions. 

The experiments were performed with suspensions of envelope-free chloroplasts isolated from spinach at 20 C temperature. The initial pH was 8.0 and did never deviate more than 0.05 units. The intensity of the actinic light was 300 Wm-2 unless otherwise stated. Changes of external pH (by glass electrode measurement) and of fluorescence (from the added indicator of Internal pH) were monitored simultaneously. Changes of the ATP concentration were obtained from the change of external pH (0.94 H consumed per ATP formed). Light-induced electron flow (in the presence of ferricyanide as electron acceptor) was obtained from the decrease of external pH (1 H liberated per electron transported). Transmembrane ApH was obtained from the fluorescence quenching of N-(l-naphtyl)ethylenediamine (NED) or from the external pH jump after a pulse of 50 pM imidazole. The H/e determination was performed as described in [1]. 

Deviation of the electrode function from the theoretical one is the other source of error. In connection with this, the well-known acidic and basic error of pH measuring glass electrodes should be mentioned. 

TABLE I. Mean deviation of pH obtained with appropriate phosphate standarized antimony electrodes from the pH measured with a glass electrode, for the same solutions containing bicarbonate and acetate. (Figures in parentheses indicate 3 solutions of bicarbonate measure with 7 electrodes, and 4 solutions of acetate measured with 8 electrodes). 

As mentioned previously, a pH glass electrode and a calomel reference electrode immersed in pH 7 buffer generate approximately zero millivolts. The standardization or calibration control on the pH meter allows any deviation of these electrodes from zero millivolts to be offset, thus providing comparison of standardization of all electrode pairs. This provides a starting point for a linear function which represents the observed potential versus the pH, and since the slope of this function is predictable, it can be represented by the function illustrated in Figure 1.6. 

If the two pairs of electrodes show a pH difference in a sample of known pH value after being standardized on the same standard pH buffer solution, then the faulty electrode can be pinpointed. If the deviation is observed in a sample of unknown value, the problem can be traced only to either reference or either glass electrode. In other words, if pairs (A) and (C) and pairs (B) and (D) each agree on the pH value, the source of difficulty would be with one of the reference electrodes. If pairs (A) and (B) and pairs (C) and (D) each agree, the source of difficulty would be with one of the glass electrodes. At this point, individual electrode tests previously described in this chapter should be initiated to aid in specific identification of the faulty electrode. 

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