complexity of conventional electroanalytical techniques results from the simultaneous occurrence of diffusion, charge transfer, and kinetics. By removing diffusion from the picture, one obtains the simplicity of homogeneous kinetic methods combined with the additional advantage of potential control. It is anticipated that these important advantages will outweigh the potential problems alluded to above.
nonhomogeneous potential on the working electrode. This effect leads to deviations from the applied potential which are larger for large currents and would cause negative deviations from theory for higher R values. In this case, however, the current decreases greatly after the initial electrolysis, to the value required to maintain the oxlred ratio as the chemical reaction proceeds. This current is generally less than 10 PA, usually on the order of 1-3 PA, values which should result in negligible error in aqueous solutions. Further evidence that ohmic losses are negligible is provided by the fact that halving the PAP concentration, resulting in smaller total current, caused no observable change in rate constants at the half-wave potential. The third source of error is more difficult to avoid, that of a slow charge transfer rate. If the chemical reaction becomes fast relative to the charge transfer rate, the oxlred ratio will not be maintained, and the result will be a drawn-out kobsd vs. Eappplot, similar to a quasi-reversible polarogram. A likely means to remove this problem is the addition of a redox mediator to the solution. If the mediator interacts rapidly with both the electrode and the couple of interest, the solution potential will be maintained at the proper value, and the oxlred ratio will be a constant. Finally, it should be noted that all these effects can cause uncertainty in the starting time for the reaction. Although not important for first-order reactions, the uncertain starting time will be more critical for higher order systems. In these cases, care must be taken to assure that the kinetic run is long relative to the time required to establish the oxlred ratio. It should be emphasized that the strengths of this thin-layer kinetic method all result from the fact that one is working with a homogeneous solution in equilibrium with an electrode. The
ACKNOWLEDGMENT The author thanks William Heineman of the University of Cincinnati for helpful discussions during this work. LITERATURE CITED (1) R. N. Adams, "Electrochemistry at Solid Electrodes", Marcel Dekker, New York, N.Y., 1969. (2) A. C. Testa, and W. Reinmuth, Anal. Chem., 32, 1512 (1960). (3) W. N. Schwarz and I. Shain, J. Phys. Chem., 69, 30 (1965). ( 4 ) R. S. Nicholson and i. Shain, Anal. Chem., 36, 706 (1964). (5) M. Hawley and S. W. Feldberg, J. Phys. Chem., 70, 3459 (1966). (6) See, for example, L. Marcoux, J. Am. Chem. Soc., 93, 537 (1971). (7) T. Kuwana. Ber. Bunsenges. Phys. Chem., 77,858 (1973). (8) H. N. Blount, J. flectroanal. Chern., 42, 271 (1973). (9) R. W. Murray, W. R . Heineman, and G. W. O'Dom, Anal. Chem., 39, 1666 (1967). (IO) W. R. Heineman, 8.J. Norris, and J. F. Goetz, Anal. Chem., 47, 79 (1975). (1 1) Reference 1, p 336. (12) J. F. Corbett, J. Chem. SOC.5, 213 (1969). (13) L. Marcoux, J. Phys. Chem., 76, 3254 (1972). (14) C. R. Christensen and F. C. Anson, Anal. Chem., 36, 495 (1964).
RECEIVEDfor review September 7,1976. Accepted November 8,1976. The work was supported in part by a Cottrell Research Grant from the Research Corporation and by a grant from the Merck Company Foundation.
Activity Coefficients and Osmotic Coefficients in Aqueous Solutions of Choline Chloride at 25 OC J. B. Macaskill, M. S. Mohan,' and Roger G. Bates" Department of Chemistry, University of Florida, Gainesville, Fla. 326 1 1
Th? mean ionic activity coefficients of choline chloride (trimethyl-p-hydroxyethylammonium chloride) have been determined in aqueous solutions at molalities up to 4 mol kg-' by measurements with an organic cation-selective electrode. The results are compared with those derived from a parallel isopiestic study of the osmotic coefflclents in these solutions at molalltles up to 7 mol kg-l. Agreement between the two methods is very good, indicating that the choline Ion-selective electrode shows near-perfect Nernstian behavior over the molality range 0.001 to 4 mol kg-I.
The choline ion is an essential component of biological systems. It is important in the metabolism of carbohydrates and nitrogen compounds and plays a role in lipid metabolism as well. The behavior of choline chloride (trimethyl-(3-hydroxyethylammonium chloride) has been studied previously by the gravimetric isopiestic technique. Fleming ( 1 ) concluded, on the basis of his experimental osmotic coefficient data, that the behavior of choline chloride in aqueous solutions Present address, D e p a r t m e n t of Chemistry, Texas versity, College Station, Texas.
A&M Uni-
is almost identical to that of ammonium chloride. This result is surprising in view of the structural similarity of the choline and tetramethylammonium ions and the established fact that the behavior of Me4NC1 in aqueous solutions is significantly different from that of NH&l(2). The osmotic coefficients (4) at a molality of 0.1 mol kg-l are 0.927 for NH4Cl and 0.910 for Me4NC1,and the difference increases with concentration. Boyd, Schwarz, and Lindenbaum ( 3 ) took note of this anomaly and made further isopiestic measurements on choline chloride. They concluded that choline chloride is indeed very similar to tetramethylammonium chloride and, in fact, that the osmotic and activity coefficients for choline chloride are even lower than those of Me4NC1. One might have expected any changes through replacement of a methyl group by a phydroxyethyl group not only to be slight but to be in the direction of increasing the osmotic and activity coefficients. Boyd et al. ( 3 )thus concluded that the (3-hydroxyethyl group must produce a sizable dipole moment in the cation. Whatever the explanation, the decreased activity of choline chloride relative to Me4NC1is apparently substantiated by conductance studies ( 4 , 5 )which show that the limiting conductivity of the choline ion is significantly lower than that of MedN+. Baum (6) has made measurements on solutions of choline and acetylcholine chlorides using an organic cation-selective
ANALYTICAL CHEMISTRY, VOL. 49, NO. 2, FEBRUARY 1977
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Table I. Electromotive Force ( E ,in mV) of the Cell: Ag;AgClIChCl(m)lCholine ISE at 25 "C.Mean Activity Coefficient (y+) of Choline Chloride m, mol kg-l 0.001
0.001858 0.003766 0.007483 0.009516 0.010942 0.016834 0.0 2938 0.04274 0.05503 0.06517 0.08714 0.10164 0.18849 0.3246 0.5200
" Calibration data: E"
E -264.5 -233.3 -198.2 -164.25 -152.55 -145.75 -124.9 -98.3 -80.65 -69.0 -61.15 -47.7 -40.95 -13.65 9.8 30.8
m , mol kg-l
Y+
0.9644" 0.9523" 0.9337" 0.9096" 0.8994" 0.8931" 0.872 0.839 0.814 0.793 0.781 0.754 0.742 0.682 0.626 0.588
0.6877 0.9706 1.1571 1.2849 1.4992 1.7103 1.9301 2.2562 2.4852 2.7133 2.937 3.269 3.500 3.863 4.084 4.265 4.377
E 41.9 56.9 64.4 69.4 76.1 81.95 87.5 95.1 100.05 104.35 108.65 114.6 118.4 124.3 127.7 130.45 132.1
Y+ 0.552 0.524 0.509 0.505 0.493 0.485 0.479 0.475 0.475 0.473 0.475 0.480 0.483 0.490 0.496 0.501 0.504
= 91.55 mV; emf slope for Choline ISE = 58.90 mV/decade.
Table 11. Molalities ( m )and Osmotic Coefficients (4) of Isopiestic Solutions of NaCl and Choline Chloride (ChCl) m (NaCl) 0.1435 0.3052 0.4149 0.7243 0.8898 1.1097 1.1301 1.2356 1.3636 1.5123 1.6160 1.7761 1.9317 2.0753 2.4681
m (ChC1)
0.1483 0.3208 0.4419 0.7867 0.9719 1.2212
1.2444 1.3678 1.5107 1.6776 1.7960 1.9828 2.1584 2.3252 2.7727
4 0.8979 0.8760 0.8641 0.8526 0.8525 0.8537 0.8539 0.8535 0.8580 0.8632 0.8663 0.8698 0.8766 0.8814 0.8999
electrode. He used both the silver-silver chloride electrode and a saturated calomel electrode as references and normalized on the results of Fleming ( 1 ) for the activity coefficient of choline chloride. In this way, he was able to show that the indicator electrode displayed Nernstian response t o the acetylcholine ion and near-Nernstian response t o the choline ion down to concentrations as low as M. As part of our continuing studies of mixed electrolytes, we have made further isopiestic measurements on solutions of choline chloride t o molalities as high as 7 mol kg-l. Activity coefficients derived from this study are compared with those obtained directly from the emf of the cell without liquid junction Ag;AgClI ChCl(m) ICholine ISE
(A)
where Ch represents choline and m is molality. Molalities as high as 4 mol kg-l were studied as a test of the response of the choline ion-selective electrode in concentrated solutions.
EXPERIMENTAL Choline chloride of reagent grade was obtained from Sigma Chemical Company. It was twice recrystallized from ethanol and dried in vacuum for a prolonged period at a temperature near 80 O C . The crystalline solid was periodically reground during the drying process and was finally dried in vacuum at 130 O C . It was analyzed gravimetrically and found t o assay 99.95 f 0.03% on the basis of its chlorine content. Solutions were prepared by appropriate dilution of a stock 210
m (NaCl)
m (ChCl)
6
2.6615 2.7638 3.0120 3.2490 3.3533 3.5000 3.5175 3.7906 3.9194 4.1287 4.3742 4.6926 5.2303 5.6222 6.1330
3.0054 3.1193 3.4034 3.6775 3.8056 3.9643 3.9984 4.3215 4.4655 4.6894 4.9747 5.3561 6.0102 6.4858 7.1240
0.9060 0.9123 0.9257 0.9383 0.9422 0.9531 0.9508 0.9652 0.9741 0.9907 1.0056 1.0233 1.0527 1.0754 1.1030
solution which had been standardized by the gravimetric AgCl procedure. The equipment for isopiestic measurements was of very simple design. It consisted of a vacuum desiccator containing up to 12 goldplated silver cups resting on a slab of copper 2.5 cm in thickness. Sodium chloride was used as a reference electrolyte, and isopiestic equilibrium was assumed to have been reached when solutions in replicate cups had the same molality within 0.1%; agreement was usually well within this limit. Glass beads were placed in the solution cups to aid agitation. The desiccator was carefully evacuated to a pressure of about 25 Torr. It was then immersed in a water bath maintained at 25 f 0.01 "C, where it was rocked gently by mechanical means to hasten the attainment of vapor equilibrium. For concentrated solutions, a period of two to three days was sufficient, but for the most dilute solutions, periods of up to one month were required. The choline ion electrode was prepared by placing the ion-selective membrane in an Orion Series 92 electrode body. The sensing membrane consisted of the exchanger p-chlorotetraphenylborate immobilized in PVC with a suitable plasticizer. The inner reference electrode was a chloridized silver wire, and the internal solution was 4 M KC1 saturated with AgC1. The silver-silver chloride electrodes of cell A were prepared by the common thermal-electrolytic procedure. Potentiometric measurements to 0.1 mV were made with a Corning Model 112 digital pH meter. Test solutions were maintained at 25 "C in a water-jacketed cell and were stirred during the measurements. The electrodes were brought t o equilibrium initially in a 0.001 m solution of choline chloride, and the molality of ChCl was then increased by stepwise addition of known weights of stock solution. The response of the ion-selectiveelectrode was very rapid, but typically 30 min was
ANALYTICAL CHEMISTRY, VOL. 49, NO. 2, FEBRUARY 1977
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~
Table 111. Parameters of Equation 3 for the Osmotic Coefficients of NaCl and ChCl at 25 "C NaCl ChCl
A lo2 B
1.46352 2.06703 10.4152 -80.652 2.17310
103 c
105 D 105 E
Std dev (6)
0.68452 2.31637 4.61580 -4.70822 -2.92897 0.0018
allowed between changes of concentration in order to assure that the Ag;AgCl electrodes had reached equilibrium. After equilibration,the cell emf was constant within 0.1 mV for a period of several hours.
RESULTS Emf Measurements. Table I summarizes the experimental results obtained over the range of molalities 0.001 to 4.4 mol kg-l. The Nernst equation for cell A is written E=E"+klogmy,
(1)
where E", k , and y i are all unknown. In the procedure of Baum ( 6 ) , the slope (k) was evaluated by normalizing on a value for the activity coefficient at a molality of 0.1 mol kg-I from the isopiestic measurements of Fleming (1). Some functional variation of y+ with ionic strength a t lower molalities was then assumed. We have attempted to minimize the arbitrary nature of this normalization step by fitting the values of y+ at 0.1 mol kg-I (y+ = 0.742 from our isopiestic data) to the equation
where I is the ionic strength and B M , Xis an interaction parameter of the BrQnsted-Guggenheim type. This expression was then used to calculate y+ over the molality range 0.01 to 0.001 mol kg-l. By this procedure, all measurements at molalities greater than 0.01 mol kg-l gave an unweighted experimental activity coefficient. By applying this approach to the appropriate results (at the six lowest molalities) in Table I, the following values of E" and k at 25 O C were obtained: E" = 91.55 f 0.13 mV k = 118.06 f 0.05 mvldecade During the course of the investigation, three similar cationic electrodes with different concentrations of internal KCl solution were assembled. Although the emf slopes of these electrodes were detectably different, they varied by no more than f0.25 mvldecade. If both constancy of slope over the range of solution concentrations studied and the absence of a time-dependent variation of E" are assumed, activity coefficients can be calculated by applying Equation 1 to the remainder of the data. The results of this calculation are shown in Table I. Isopiestic Measurements. The results of the isopiestic measurements are given in Table I1 as molalities of the isopiestic solutions together with the derived values of the osmotic coefficient $. The osmotic coefficient of the sodium chloride reference solutions was calculated from an expression of the Lietzke-Stoughton (7) type
1
+ A1112- 2 ln(1 + AI1/*) + BI + + 0 1 3 + E I ~ (3) c12
where A, B, C, D, and E are adjustable parameters and the Debye-Huckel slope S has the value 1.17202. The appropriate values of A , B , C, D , and E are listed in Table 111. The exper-
m Flgure 1. Osmotic coefficient of choline chloride solidions as a function of molality. (0)present work: ( 0 )smoothed data of Boyd et al. (3)
0.
t Yf
0
0
m Flgure 2. Mean activity coefficient of choline chloride as a function of molality at 25 "C (0)present work (isopiestic);( 0 )present work of Boyd et al. (3)
(emf). Lower curve from data
imental osmotic coefficients for choline chloride were then fitted to the same expression; the best values of the parameters for this system are also given in Table 111. Activity coefficients were then calculated directly with the aid of the equation (7) -SI1/2 2BI -3C 1 2 + -4D I 3 + -5E 1 4 ( 4 ) In yi = 1 AI112 2 3 2 The present results are in general agreement with those of Boyd et al. ( 3 ) ,supporting their finding that qbChCl< $ M ~ ~ N < $ N H ~ C I .In Figure 1, both sets of osmotic coefficients are plotted as a function of molality, and it can be seen that there
+
+
+
ANALYTICAL CHEMISTRY, VOL. 49, NO. 2, FEBRUARY 1977
211
C ~
is nonetheless a definite disparity between the two sets of data, with a maximum difference of about 0.01 in the osmotic coefficient near a molality of 1mol kg-l. We have repeated our measurements with different preparations of choline chloride but have been unable to reduce materially this systematic discrepancy.
DISCUSSION In Figure 2, the mean ionic activity coefficients of choline chloride derived from the three separate studies are plotted as a function of molality. It is evident that the isopiestic data are in good agreement with the results of our emf measurements. In the molality range 1to 3 mol kg-l, there is a systematic difference of up to 0.007 in yh,but a single curve fits the data to within f0.004 in y&.It is difficult to favor either set of results. With the emf measurements, one might expect an apparent reduced activity coefficient at higher molalities due to the progressive formation of chloro complexes as the solubility of AgCl is enhanced by the increased concentration of chloride ion. Nevertheless, the two sets of data are in closer agreement at high molalities, where the effect should be most pronounced, than at low and moderate concentrations. The isopiestic measurements are of adequate precision, yet there is always an exaggeration of experimental uncertainties when transforming $ into y&,because of the form of the expression relating these two quantities:
This point is evident in the appreciable discrepancy between the activity coefficients derived from the two sets of osmotic coefficient data. The difference in $ is at most 0.01, yet the values of 7%differ by as much as 0.025. This is not an artifact
of the method used by Boyd et al. (3) to derive yh from 4. We have recalculated their data using the equations of Pitzer (8) and obtain activity coefficients in good agreement with those they reported. In terms of selectivity, stability, response time, and approach to the theoretical Nernst slope, the choline ion-selective electrode compares favorably with existing commercial ion-selective electrodes. In addition to its application in clinical and physiological situations, it has uses in fundamental studies in solution chemistry. In the investigation of salt systems such as NaCl/ChCl, electrodes reversible to Na+ and Ch+ permit the mean ionic activity coefficients of both components to be measured and thus provide a comparison with the predictions of current treatments of mixed electrolytes.
ACKNOWLEDGMENT The authors are indebted to G. Baum for supplying the choline-selective membranes and to Alan Pethybridge for assistance in the computation of y* from 4.
LITERATURE CITED (1) R. Fleming, J. Chem. Soc., 3100 (1961). (2) S. Lindenbaum and G. E. Boyd, J. Phys. Chem., 68, 91 1 (1964). (3) G. E. Boyd, A. Schwarz, and S. Lindenbaum, J. Phys. Chem., 70, 621 (1966). (4) R. Fleming, J. Chem. Soc., 4914 (1960). (5) H. 0. Spivey and F. M. Snell, J. Phys. Chem., 68, 2126 (1964). (6) G. Baum, Anal. Lett., 3, 105 (1970). (7) M. H. Lietzke and R. W. Stoughton, J. Phys. Chem., 65, 508 (1961). (8) K. S.Pitzer, J. Phys. Chem., 77, 268 (1973).
RECEIVEDfor review October 8, 1976. Accepted November 1, 1976. This work was supported in part by the National Science Foundation under Grant CHE73-05019 A02.
Monoionic Liquid Membrane Potential in Non-Isothermal Systems Giancarlo Scibona, Mauro Magini, Bernardlno Scuppa, Anna Castagnola, and Claudio Fabiani” Material Division, C.N.E.N., CSKCasaccia, Rome, /taw
Liquid membranes formed by alkylammonium salts in suitable organic solvents behave as specific electrodes when interposed between two aqueous electrolyte solutlons. A thermodynamic analysis of matter and energy transport through the liquid ion exchange electrodes has shown ( 4 ) that for monoionlc system the thermal membrane (electrode) potential has to depend on the nature of the counterlon, on the aqueous actlvity of the counterion and on the nature of the solvent phase. The experimental thermal membrane potentials (tetraheptylammonium salts In benzene, monochlorobenzene, o-dichlorobenzene and nitrobenzene) studled in this paper fully agree wlthlhese predictions and represent a contribution to the knowledge of the temperature dependence of speclflc liquid electrodes of Interest to analytical chemistry.
Solutions of alkylammonium salts in low dielectric constant solvents behave as liquid anion exchange electrodes when interposed between two aqueous solutions of suitable composition ( I ) . A mechanistic theory based on the coupling of ion and ion pair diffusion processes with ion exchange at the water-membrane interfaces, has been developed ( 2 ) to de212
scribe the electrochemical behavior of these membranes in isothermal conditions. Any investigation of the thermal membrane potential of liquid ion exchange membranes requires the understanding of the phenomena associated with the transport of matter and energy through the membranes. In the case of solid ion exchange membranes, the electrical membrane potential in non-isothermal conditions is given by ( 3 ) -A$ = ( 2 t + - l)(RT/F) In ( a * ” / u * ’ ) (2t+ - l ) ( R A T / F )In a*’’ (t+a+ t-a-)AT
+
+
+
(1)
with CY+= TJ - s + / F and CY- = TJ - s-/F. The solvent contributions have been omitted from Equation 1. All quantities, but TJ and t ( k ) ,refer to the external aqueous phases: phases(‘) at temperature T and phase(”) at temperature T AT; u* is the mean aqueous electrolyte activity; s+, s- and t+,t - are the cation and the anion aqueous entropies and membrane transport numbers, respectively; TJ is assumed constant. In the absence of electrolyte uptake and then with co-ion exclusion, as is the case for many liquid ion exchange membranes, the membrane behaves as highly permselective to the counterions. Under these conditions the thermal potential
ANALYTICAL CHEMISTRY, VOL. 49, NO. 2, FEBRUARY 1977
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