Liquid junction potentials values

So far, a cell containing a single electrolyte solution has been considered (a galvanic cell without transport). When the two electrodes of the cell are immersed into different electrolyte solutions in the same solvent, separated by a liquid junction (see Section 2.5.3), this system is termed a galvanic cell with transport. The relationship for the EMF of this type of a cell is based on a balance of the Galvani potential differences. This approach yields a result similar to that obtained in the calculation of the EMF of a cell without transport, plus the liquid junction potential value A0L. Thus Eq. (3.1.66) assumes the form... 

Table 2.2. Liquid-junction potential values all values at 25 °Caccording to (2.6.13) compared with experiment [22].

As a result of a variable liquid-junction potential, the measured pH may be expected to differ seriously from the determined from cells without a liquid junction in solutions of high acidity or high alkalinity. Merely to affirm the proper functioning of the glass electrode at the extreme ends of the pH scale, two secondary standards are included in Table 8.14. In addition, values for a 0.1 m solution of HCl are given to extend the pH scale up to 275°C [see R. S. Greeley, Anal. Chem. 32 1717 (I960)] ... 

Since the small interfacial potentials at the junctions of the electrodes and copper leads are equal and opposite, they cancel out, and if the liquid junction potential is assumed to be small, or is reduced to a negligible value by using a salt bridge, then equation 20.216 reduces to... 

Standard potentials Ee are evaluated with full regard to activity effects and with all ions present in simple form they are really limiting or ideal values and are rarely observed in a potentiometric measurement. In practice, the solutions may be quite concentrated and frequently contain other electrolytes under these conditions the activities of the pertinent species are much smaller than the concentrations, and consequently the use of the latter may lead to unreliable conclusions. Also, the actual active species present (see example below) may differ from those to which the ideal standard potentials apply. For these reasons formal potentials have been proposed to supplement standard potentials. The formal potential is the potential observed experimentally in a solution containing one mole each of the oxidised and reduced substances together with other specified substances at specified concentrations. It is found that formal potentials vary appreciably, for example, with the nature and concentration of the acid that is present. The formal potential incorporates in one value the effects resulting from variation of activity coefficients with ionic strength, acid-base dissociation, complexation, liquid-junction potentials, etc., and thus has a real practical value. Formal potentials do not have the theoretical significance of standard potentials, but they are observed values in actual potentiometric measurements. In dilute solutions they usually obey the Nernst equation fairly closely in the form ... 

An element of uncertainty is introduced into the e.m.f. measurement by the liquid junction potential which is established at the interface between the two solutions, one pertaining to the reference electrode and the other to the indicator electrode. This liquid junction potential can be largely eliminated, however, if one solution contains a high concentration of potassium chloride or of ammonium nitrate, electrolytes in which the ionic conductivities of the cation and the anion have very similar values. 

In view of the problems referred to above in connection with direct potentiometry, much attention has been directed to the procedure of potentio-metric titration as an analytical method. As the name implies, it is a titrimetric procedure in which potentiometric measurements are carried out in order to fix the end point. In this procedure we are concerned with changes in electrode potential rather than in an accurate value for the electrode potential with a given solution, and under these circumstances the effect of the liquid junction potential may be ignored. In such a titration, the change in cell e.m.f. occurs most rapidly in the neighbourhood of the end point, and as will be explained later (Section 15.18), various methods can be used to ascertain the point at which the rate of potential change is at a maximum this is at the end point of the titration. 

The most widely used reference electrode, due to its ease of preparation and constancy of potential, is the calomel electrode. A calomel half-cell is one in which mercury and calomel [mercury(I) chloride] are covered with potassium chloride solution of definite concentration this may be 0.1 M, 1M, or saturated. These electrodes are referred to as the decimolar, the molar and the saturated calomel electrode (S.C.E.) and have the potentials, relative to the standard hydrogen electrode at 25 °C, of 0.3358,0.2824 and 0.2444 volt. Of these electrodes the S.C.E. is most commonly used, largely because of the suppressive effect of saturated potassium chloride solution on liquid junction potentials. However, this electrode suffers from the drawback that its potential varies rapidly with alteration in temperature owing to changes in the solubility of potassium chloride, and restoration of a stable potential may be slow owing to the disturbance of the calomel-potassium chloride equilibrium. The potentials of the decimolar and molar electrodes are less affected by change in temperature and are to be preferred in cases where accurate values of electrode potentials are required. The electrode reaction is... 

A typical set of experimental data290a,290b is shown in Fig. 11. All measurements converge to the value measured by Grahame.286 At present, the of Hg in water can be confidently indicated5 as -0.433 0.001 V (SCE), i.e., -0.192 0.001 V (SHE). The residual uncertainty is related to the unknown liquid junction potential at the boundary with the SCE, which is customarily used as a reference electrode. The temperature coefficient of of the Hg/H20 interface has been measured and its significance discussed.7,106,1 8,291... 

Japaridze et al.m 323 have studied the interface between Hg and a number of vicinal and nonvicinal diols such as 1,2-, 1,3-, 2,3- and 1,4-butanediol (BD), ethanediol (ED), and 1,3-propanediol. KF and LiC104 were used as surface-inactive electrolytes. The potential of zero charge was measured by the capacitance method against an SCE in water without correction for the liquid junction potential at the solvent/H20 contact (such a potential drop is estimated to be in the range of 20 to 30 mV). The potential of the capacitance minimum was found to be independent of the electrolyte concentration while capacitance decreased with dilution. Therefore, Emin was taken to measure E . These values are reported in Table 4. 

Difference with respect to value calculated from accepted standard potential and activity coefficient is attributed to liquid junction potential. 

The two equations show that the nearer the cationic transport is to 0.5, the smaller is the liquid junction potential (other factors being unchanged). Among common electrolytes one of the highest numerical values of the factor (2 t+ - I) is given by hydrochloric acid, at 0.65. Hence a potential difference of about 39 mV develops at 25 °C across the junction between 0.001 N and 0.01 N hydrochloric acid. In the case of potassium chloride solution,... 

When eliminating the liquid junction potential by one of the methods described in Section 2.5.3, we obtain a concentration cell without transport. The value of its EMF is given simply by the difference between the two electrode potentials. More exactly than by the described elimination of the liquid junction potential, a concentration cell without transport can be obtained by using amalgam electrodes or electrodes of the second kind. 

It would appear from Eq. (3.2.8) that the pH, i.e. the activity of a single type of ion, can be measured exactly. This is not, in reality, true even if the liquid junction potential is eliminated the value of Eref must be known. This value is always determined by assuming that the activity coefficients depend only on the overall ionic strength and not on the ionic species. Thus the mean activities and mean activity coefficients of the electrolyte must be employed. The use of this assumption in the determination of the value of Eref will, of course, also affect the pH value found from Eq. (3.2.8). Thus, the potentiometric determination of the pH is more difficult than would appear at first glance and will be considered in the special Section 3.3.2. 

A liquid junction potential E-f forms when the two half-cells of a cell contain different electrolyte solutions. The magnitude of Ej depends on the concentrations (strictly, the activities) of the constituent ions in the cell, the charges of each moving ion, and on the relative rates of ionic movement across the membrane. We record a constant value of j because equilibrium forms within a few milliseconds of the two half-cells adjoining across the membrane. 

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. 

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. 

Theoretical cell potential The algebraic sum of the individual redox potentials of an electrochemical cell at zero current, i.e. emf = Epositive electrode - negative electrode- In practice, when Current flows in a cell, a liquid junction potential is present, and the cell potential is larger than this theoretical value. 

An acceptable method quite frequently used in practice depends on the cell whose EMF is being measured having a liquid junction with a constant potential value. Such a situation is attained in the determination of the activity of fluoride ions, by adding a constant amount of quite concentrated buffer, for example TISAB, to the studied solution this buffer also fulfills other functions in the analysis (see p. 146). Then the liquid junction potential is a function of the composition of the reference electrode electrolyte and of the buffer composition alone, and not of the concentrations of the other components of the studied solution. 

Another less precise but frequently used method employs a liquid bridge between the analysed solution and the reference electrode solution. This bridge is usually filled with a saturated or 3.5 m KCl solution. If the reference electrode is a saturated calomel electrode, no further liquid bridge is necessary. Use of this bridge is based on the fact that the mobilities of potassium and chloride ions are about the same so that, as follows from the Henderson equation, the liquid-junction potential with a dilute solution on the other side has a very low value. Only when the saturated KCl solution is in contact with a very concentrated electrolyte solution with very different cation and anion mobilities does the liquid junction potential attain larger values [2] for the liquid junction 3.5 M KCl II1 M NaOH, A0z, = 10.5 mV. 

The calibration of ISEs using the tabulated activity values is especially simple when the dependence of the ISE potential on the determinand activity is Nemstian (3.1.5). If the liquid-junction potential is negligible or constant, the determinand activity flj+(X) can be found, using the standard activity [Pg.80]

The ionic strength of the solution influences the activity coefficients of the species present and the values of the liquid-junction potentials in the system and thus must be held constant in all related measurements, within a range of about 0.1 to 2m. 

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. 

Examination of the behaviour of a dilute solution of the substrate at a small electrode is a preliminary step towards electrochemical transformation of an organic compound. The electrode potential is swept in a linear fashion and the current recorded. This experiment shows the potential range where the substrate is electroactive and information about the mechanism of the electrochemical process can be deduced from the shape of the voltammetric response curve [44]. Substrate concentrations of the order of 10 molar are used with electrodes of area 0.2 cm or less and a supporting electrolyte concentration around 0.1 molar. As the electrode potential is swept through the electroactive region, a current response of the order of microamperes is seen. The response rises and eventually reaches a maximum value. At such low substrate concentration, the rate of the surface electron transfer process eventually becomes limited by the rate of diffusion of substrate towards the electrode. The counter electrode is placed in the same reaction vessel. At these low concentrations, products formed at the counter electrode do not interfere with the working electrode process. The potential of the working electrode is controlled relative to a reference electrode. For most work, even in aprotic solvents, the reference electrode is the aqueous saturated calomel electrode. Quoted reaction potentials then include the liquid junction potential. A reference electrode, which uses the same solvent as the main electrochemical cell, is used when mechanistic conclusions are to be drawn from the experimental results. 

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. 

Undoubtedly, the mercury/aqueous solution interface, was in the past, the most intensively studied interface, which was reflected in a large number of original and review papers devoted to its description, for example. Ref. 1, and in the more recent work by Trasatti and Lust [2] on the potentials of zero charge. It is noteworthy that in view of numerous measurements of the double-layer capacitance at mercury brought in contact with NaF and Na2S04 solutions, the classical theory of Grahame [3] stiU holds [2]. According to Trasatti [4], the most reliable PZC value for Hg/H20 interface in the absence of specific adsorption equals to —0.433 0.001 V versus saturated calomel electrode, (SCE) residual uncertainty arises mainly from the unknown liquid junction potential at the electrolyte solution/SCE reference electrode boundary. 

Recently, Fuchs etal. [15], using the streaming mercury electrode and applying the Henderson equation, have determined the pzc value in the solutions of tetraethy-lammonium perchlorate in DMSO as —0.515 0.001 V (versus Ag/0.01 M Ag+ (DMSO) reference electrode). This value was corrected for the liquid junction potential and was independent of tetraethyl ammonium perchlorate (TEAR) concentration within the range 0.02 to 0.75 M. Using the same methodology, KiSova et al. 

S. Mine, J. Jastrzebska, Rocz. Chem. 1954, 28, 519-520. In reporting the E20 value for copper in various methanol-water mixtures, these authors did not specify whether percent methanol referred to weight, volume, or mole per cent. Although the potential values are claimed to be referenced to aqueous SHE, the actual internal reference electrode used was not specified and no mention was made of a correction for the liquid junction potential. Since the actual liquid junction potential for aqueous reference electrodes in methanol-water mixtures was not made until a later date [ j —0.152 V for 80% methanol (w/w) M. Alfenaar, C. L. deLigny, Reel. Trav. Chim. Pays-Bas... 

Most pH meters will display the electrode response (slope) as a percentage of the theoretical value (59.16 mV/pH unit at 25°C). The electrode response should not be less than 95% or more than 105% of the theoretical value at a given buffer temperature. Contamination of the electrode or changes in the liquid junction potential will typically lower the electrode response to below 95%. Some common procedures to regenerate the electrode are listed below. Replacement of... 

The potential values contain the liquid junction potentials at 0.1 M Et4NClO4(AN)/0.1 M Et4NCl04(S). 

If two electrolyte solutions that are of different concentrations but in the same solvent contact each other at a junction, ion transfers occur across the junction (Fig. 6.3). If the rate of transfer of the cation differs from that of the anion, a charge separation occurs at the junction and a potential difference is generated. The potential difference tends to retard the ion of higher rate and accelerate the ion of lower rate. Eventually, the rates of both ions are balanced and the potential difference reaches a constant value. This potential difference is called the liquid junction potential (LJP) [10]. As for the LJP between aqueous solutions, the LJP between non-aqueous solutions can be estimated using the Henderson equation. Generally the LJP, Lj-, at the junction Ci MX(s) c2 NY(s) can be expressed by Eq. (6.1) ... 

Polarographic reductions have been studied for many kinds of metal ions and in a variety of lion-aqueous solvents. Large amounts of data on half-wave potentials are available and have been compiled in some books and reviews [18]. However, many of the old data were obtained using different reference electrodes or aqueous SCE, for which the problem of the liquid junction potential exists (Section 6.1.2). Gritzner [19] compiled half-wave potentials for metal ions as values referred to the BCr+/BCr system, which was recommended by IUPAC. Some of these are listed in Table 8.2. The potential of the BCr+/BCr system is not seriously affected either by the presence of water and other impurities or by differences in experimental conditions. Thus, although the determination of the half-wave poten-... 

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