Error junction potential

The temperature compensator on a pH meter varies the instrument definition of a pH unit from 54.20 mV at 0°C to perhaps 66.10 mV at 60°C. This permits one to measure the pH of the sample (and reference buffer standard) at its actual temperature and thus avoid error due to dissociation equilibria and to junction potentials which have significant temperature coefficients. 

Reference electrodes for non-aqueous solvents are always troublesome because the necessary salt bridge may add considerable errors by undefined junction potentials. Leakage of components of the reference compartment, water in particular, into the working electrode compartment is a further problem. Whenever electrochemical cells of very small dimensions have to be designed, the construction of a suitable reference electrode system may be very difficult. Thus, an ideal reference electrode would be a simple wire introduced into the test cell. The usefulness of redox modified electrodes as reference electrodes in this respect has been studied in some detail... 

Table IV lists the redox potentials of conjugated ferrocene oligomers (mainly dimers with a single bridge). Potential values are denoted against different reference electrodes as given in the references. The values can be primarily compared using the relationship mentioned in the footnote of the table, although care should be taken with some errors derived from junction potentials which depend on experimental conditions. There have been several reports on the quantitative estimation of structural factors affecting internuclear electron delocalization.

Although the Henderson formula depends on a number of simplifications and also employs ion concentrations rather than activities, it can nonetheless be used for estimation of the liquid junction potential up to moderate electrolyte concentrations, apparently as a result of compensation of errors. 

Liquid junction potentials are rarely large, so a value of E as large as 0.1 V should be regarded as exceptional. Nevertheless, junction potentials of 30 mV are common and a major cause of experimental error, in part because they are difficult to quantify, but also because they can be quite irreproducible. 

To leam about how a liquid junction potential, j, arises, and appreciate how it can lead to significant errors in a calculation which uses potentio-metric data. 

An electrode potential Eaq+.aq is measured as 0.670 V. In fact, this value is too high because the value of Eaq+.aq also incorporates a liquid junction potential of 22 mV. Calculate two values of a(Ag ), first by assuming that Eaq+.aq is accurate, and secondly by taking account of E. What is the error in a(Ag ) caused by the liquid junction potential Take E + = 0.799 V at 298 K. 

When considering potentiometric errors, it is necessary to appreciate how a liquid junction potential, Ej, arises, and appreciate how such potentials can lead to significant errors in a calculation. In addition, we saw how the IR drop can affect a potentiometric measurement (and described how to overcome this). Finally, we discussed the ways that potentiometric measurements are prone to errors caused by both current passage through the cell, and by the nature of the mathematical functions with which the Nemst equation is formulated. 

The results obtained with ISEs have been compared several times with those of other methods. When the determination of calcium using the Orion SS-20 analyser was tested, it was found that the results in heparinized whole blood and serum were sufficiently precise and subject to negligible interference from K and Mg ([82]), but that it is necessary to correct for the sodium error, as the ionic strength is adjusted with a sodium salt [82], and that a systematic error appears in the presence of colloids and cells due to complexa-tion and variations in the liquid-junction potential [76]. Determination of sodium and potassium with ISEs is comparable with flame photometric estimation [39, 113, 116] or is even more precise [165], but the values obtained with ISEs in serum are somewhat higher than those from flame photometry and most others methods [3, 25, 27, 113, 116]. This phenomenon is called pseudohyponatremia. It is caused by the fact that the samples are not diluted in ISE measurement, whereas in other methods dilution occurs before and during the measurement. On dilution, part of the water in serum is replaced by lipids and partially soluble serum proteins in samples with abnormally increased level of lipids and/or proteins. 

The experimental apparatus consists essentially of a narrow vertical glass tube down the inner surface of which one liquid is made to flow, the other liquid emerges from a fine glass tip in the form of a narrow jet down the axis of the tube. The two solutions are connected with calomel electrodes employing potassium chloride or nitrate as junction liquids. The E.M.F. of the cell is measured by means of a sensitive quadrant electrometer. The greatest source of error in the method is the elimination of or the calculation of the exact values of the liquid-liquid junction potentials in the system. For electrolytes which are not very capillary active, the possible error may amount to as much as fifty per cent, of the observed E.M.F. 

The difference in the conductivity of the calibration buffers and sample can cause a very large error on the sample measurement, due to junction potentials in different environments. Solid samples should be dissolved in purified water. It is necessary that the water be carbon dioxide-free. The presence of dissolved carbon dioxide will cause significant bias in the measurement of samples with low buffering capacity. For pH measurements with an accuracy of 0.01 to 0.1 pH unit, the limiting factor is often the electrochemical system (i.e., the characteristics of the electrodes and the solution in which they are immersed). 

Box 15-1 Systematic Error in Rainwater pH Measurement The Effect of Junction Potential... 

Accuracy and Interpretation of Measured pH Values. To define the pH scale and pertnil the calibration of pH measurement systems, a scries of reference buffer solutions have been certified hy the U.S. National Institute of Standards and Technology iNIST). The acidity function which is the experimental basis for the assignment of pH. is reproducible within about O.IKl.I pH unit from It) to 40T. However, errors in the standard potential of the cell, in the composition of the buffer materials, and in the preparation of the solutions may raise the uncertainty to 0 005 pH unit. The accuracy of ihe practical scale may he furthei reduced to (I.Ot)X-(l.(ll pH unit as a result of variations in the liquid-junction potential. 

Experience shows that the activity coefficients on this scale stay near unity (usually within experimental error) as long as the concentrations of the reactants are kept low, say less than 10% of the concentrations of the medium ions. The activity ( concentration) of several ions, notably H+, can be determined conveniently and accurately by means of e.m.f. methods, either with or without a liquid junction. In the latter case the liquid junction potential is small (mainly a function of [H+] ) and easily corrected for (3). The equilibrium constant for any reaction, on the ionic medium scale, may then be defined as the limiting value for the concentration quotient ... 

The list of error sources continues, just to mention a few the ionic strength of the sample, the liquid-junction and residual liquid-junction potentials, temperature effects, instabilities in the galvanic cell, carryover effects, improper use of available corrections (e.g., for pH-adjusted ionized calcium or magnesium). An error analysis goes beyond the limited scope of this paper more details are presented elsewhere [10]. 

Using KOH or NaOH with pH in the range 13 to 14 as the inner electrolyte gives rise to a change in the liquid junction potential error of about 30 to 35 mV for one pH unit change in the outer test solution. Consequently, when pH in a high pH test solution gradually decreases due to carbonation, a reference electrode of this kind will drift . 

When specific ion electrodes are employed, an error of less than 1 to 2% of the activity of the ion is desirable. This requires that the potential measurement have an error of less than 0.1 to 0.2 mV the reproducibility of the junction potential becomes a critical factor in the measurement. The sleeve-type junction... 

However, there is convincing experimental evidence that this can introduce substantial errors because the residual junction potential errors are not negligible.32 The reference electrodes should be connected to one another through die sample solution in the cell ... 

The logarithmic response of ISEs can cause major accuracy problems. Very small uncertainties in the measured cell potential can thus cause large errors. (Recall that an uncertainty of 1 mV corresponds to a relative error of 4% in the concentration of a monovalent ion.) Since potential measurements are seldom better than 0.1 mV uncertainty, best measurements of monovalent ions are limited to about 0.4% relative concentration error. In many practical situations, the error is significantly larger. The main source of error in potentio-metric measurements is actually not the ISE, but rather changes in the reference electrode junction potential, namely, the potential difference generated between the reference electrolyte and sample solution. The junction potential is caused by an unequal distribution of anions and cations across the boundary between two dissimilar electrolyte solutions (which results in ion movement at different rates). When the two solutions differ only in the electrolyte concentration, such liquid junction potential is proportional to the difference in transference numbers of the positive and negative ions and to the log of the ratio of the ions on both sides of the junction ... 

When the RTIL contained in the RE is different from the target RTIL, an experimental error can be generated because of the junction potential between these RTILs. However, the use of an internal reference such as ferrocene can remove the junction potential. This is because the redox potential measured with the RE contains the same junction potential. 

Morse (21) recently reported near-equilibrium pH-stat rates in sea water at 25°C and 10 atm CO2. His rates are all less than 8 mg cm yt . Table III compares our calculations (using equation 15 and assuming surface PCO2 is equal to the bulk fluid value) of a set of Morse s rates near equilibrium, and shows generally poor agreement. In calculation of Q and aH (s), no correction for liquid junction potential error in measured pH was necessary, as in our earlier calculations for pseudo-sea water. 

In recent years care has been taken to eliminate, or reduce, as far as possible the sources of error in the evaluation of standard oxidation-reduction potentials highly dissociated salts, such as perchlorates, are employed wherever possible, and corrections are applied for hydrolysis if it occurs. The cells are made up so as to have liquid junction potentials whose values are small and which can be determined if necessary, and the results are extrapolated to infinite dilution to avoid activity corrections. One type of procedure adopted is illustrated by the case described below. ... 

---