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2,4-dinitrobenzenesulEonicacid was used in place of p-toluenesulfonic acid. The potentiometric titration of imidazole in acetic acid with 2,4-dinitrobenzenesulfonic acid, which is typical of the compounds listed in Table IV, is illustrated in Figure 3 and compared to t’he HClO, titrant. The only difference observed between the two titrants is when the sulfonate salt precipitated. In these cases, usually a n aniline derivative, a larger potential break was found for the sulfonic acid titrant. The enhancement does not’ indicate the sulfonic acid to be a stronger acid than HClO, but) rather is a result of the added driving force of precipitation. This property has been observed before (5). The wide variety of compounds listed in Tahlc IV which can be routinely anal;\-zedby HC104titration were chosen in order to $how that 2,S-dinitrobenzenesulfonic arid and HCIOi are comparable in their applications. For esample, weak (p-nitroaniline) as n-ell as strong ( I ,3-dil,henyl~uanidi:ie)ba5ee and basic coml)ounds of a variety of structures can be titrated in a similar fashion with the sulfonic arid titrant. : I variety of solvent. can be u w l . The salt, KC1, which i. also characteristic of a n amine hydrochloride such a; tetracaine hydrochloride was analyzed by the addition
of mercuric acetate (5) and titration of the acetate ion which is formed. Interestingly, the titration of strychnine sulfate, which involves the conversion of sulfate to bisulfate, illustrates the strongly acidic nature of 2,4-dinitrobenzenesulfonic acid. In several instances amines as received were used, and the purity of the amines was determined by HC104 titration. As can be seen in Table IV, good stoichiometry was obtained. In addition to the properties already mentioned, 2,4-dinitrobenzenesulfonic acid is a solid and can be used in any solvent in which it is soluble, permitting a one-solvent system. Water, however, is not completely eliminated, since the sulfonic acid occurs under normal conditions as a dihydrate. Calculation shows that 0.36 and 0.25% water would be present for 0.1N solution if 2,4dinitrobenzenesulfonic acid dihydrate and 72% HC104,respectively, were used. The effect of adding increasing amounts of water to a n amine titration was the same for both acidic titrants. ACKNOWLEDGMENT
The authors thank American Cyanamid, Pfister Chemical, and Antara Chemicals for d a t a sheets and generous samples of several of the compounds.
LIlERAlURE CllED
(1) Bruss, D. B., Wyld, G. B. A., ANAL. CHEM.29, 232 (1957). (2) , , Carr. M. H.. Brown. H. P.. J . Am. Chem.’Soc. 69,’1170 (1947). ’ (3) Caso, M. hl., Cefola. hl., Anal. Chim. Acta 21, 374 (1959). (4) Critchfield, F. E., “Organic Func-
tional Group Analysis,” Macmillan, New York, 1963. ( 5 ) Fritz, J. S. Hammond, G. S., “Quantitative Orginic Analysis,” Wiley, New York, 1957. (6) Hall, H. K., Jr., J . Phys. Chem. 60, 63 (1956). (7) Kucharsky, J., Safarik, L., “Titrations in Nonaaueous Solvents.” Elsevier. New York, 19’65. (8) Lane, E. S., Talanta 8 , 849 (1961). (9) Paul, R. C., Pahil, S. S., Anal. Chim. Acta 30, 466 (1964). (10) Paul, R. C., Pahil, S. S., Malhotra, K. C.. T’ashisht. S. K.. J. Sci. Ind. Res. 21B, 41 (1962).‘ (11) Paul, R. C., Vashisht, S. K., Malhotra, K. C., Pahil, S. S., ANAL.CHEM.34, 820 (1962). (12) Smith, T. L., Elliott, J. H., J . Am. Chem. SOC.75. 3566 11953). (13) van der Heijde, H. B.,‘ Anal. Chim. Acta 17, 512 (1957). (14) van der Heijde, H. B., Dahmen, E. A. hl. F., Zbid., 16, 378 (1957). RECEIVEDfor review March 2, 1966. Accepted April 29, 1966. First Midwest Regional American Chemical Society Meeting, Kansas City, Mo., Xovember 4-5, 1965. Financial assistance came from a grant (GM 123106-01) from the National Institutes of Health and a Du Pont Fellowship (1965-1966) for one of the authors (JB).
Analytical Study of an Iodide-Sensitive Membrane Electrode G. A. RECHNITZ,’ M. R.
KRESZ, and S. B. ZAMOCHNICK2
Department o f Chemistry and Analytical Chemistry Center, University of Pennsylvania, Philadelphia, Pa.
b Evaluation of a precipitate-impregnated membrane electrode has revealed sensitivity, selectivity, and other response characteristics well suited to the analytical utilization of such electrodes. It has been demonstrated that iodide ion can be determined by direct potentiometry at concentration 1 O-’M in aqueous levels as low as 5 solutions with relatively little interference from chloride, bromide, and other common anions. In dilute solutions, the electrode response is independent of the nature of the cations present. Excellent results were obtained in potentiometric titrations of iodide with Ag+ using the membrane electrode as an indicator electrode. The response rates of such electrodes suggest that continuous monitoring of some changing systems may also b e feasible.
x
R
in the development of ion-selective electrodes for alkali metal and other cations (8) have created renewed interest in the possibility of devising electrode systems having selective response to other classes of common ions. Indeed, it has been suggested (1) that glass electrodes with response to anions should be feasible, but no concrete results along these lines have been obtained, as yet. Most promising among recent efforts directed toLvard the development of anion-selective electrodes has been the work of Pungor et al. (6, 7 ) , who described the preparation and some preliminary evaluation of several precipitate-impregnated membrane electrodes. These membranes consist of fine particles of a sparingly soluble precipitatee.g., BaS04 or AgI-immobilized in a polymerized silicone rubber matrix. ECEXT SUCCESSES
7 9 7 04
When fabricated into suitable form, the resulting electrodes display favorable mechanical properties and good chemical durability. The concepts underlying the design of such membrane electrodes are by no means new and have recently been comprehensively reviewed by Lakshminarayanaiha (4). I n 1958, Fischer and Babcock ( 2 ) published a n excellent analytical evaluation of electrodes consisting of BaSOAmpregnated paraffin and pointed out the possible application of such electrodes to analytical potentiometry, although the actual results obtained were not encouraging. The present study was undertaken in order t o provide some quantitative inAlfred P. Sloan Fellow. Present address: Department of Chemistry, Cornell University, Ithaca, N. Y. VOL. 38,
NO. 8,
JULY 1966
973
Filling Solution (
id3!
’
Schematic of iodide membrane electrode
formation regarding the analytical usefulness of the AgI-impregnated silicone rubber membrane electrode under laboratory conditions and to suggest some possible areas of application of such electrodes to chemical measurements in solution. EXPERIMENTAL
Apparatus. All potentiometric measurements were made using a Beckman Model 76 p H meter on the expanded scale. E.ni.f. time curves of the dynamic response measurements were obtained by displaying the output signal of the p H meter on a Photovolt Model 44 potentiometric recorder a t variable chart speeds. Indicator electrodes employed in this study were the Pungor type (&I in silicone rubber matrix) and were used in conjunction with Beckman 39170 fiber-junction saturated calomel reference electrodes. According to Pungor et al. ( 7 ) , the membranes are best prepared by the cold, catalyzed polymerization of silicone rubber monomer mixed with the silver iodide precipitate. The most favorable grain size for the precipitate is 5 to 10 microns; particles of this size are apparently obtained when the precipitation is carried in the presence of organic additives to retard secondary nucleation. The thickness of the membranes used in electrode construction is approximately 0.5 millimeter. Untreated membrane electrodes were obtained from National Instrument Laboratories, Inc., Rockville, Md., and were preconditioned for the purposes of this investigation by extensive (48-hour minimum) soaking in 10-2;12 samples of the appropriate test solution after careful selection for physical characteridcs and potentiometric stability. .ipproximately 50y0 of the electrodes had t o be discarded because of leakage of excessive drift (>2 mv./hr.). Figure 1 is a schematic representation of the final membrane electrode. Calibration and response curves were constructed from data taken a t 25 i 0.5’ C. Reagents. Chemicals used were of reagent grade except for barium iodide and cerous iodide, which were of purified grade. All solutions were prepared from water which had been both deionized and distilled; no residual quantities of halides could be detected in the water used. Stock 974
Agi(*50wf. % ) in Silicone Rubber
KI)
Figure 1 .
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/
ANALYTICAL CHEMISTRY
solutions 10-lLlfin iodide were prepared by weight from KI, BaI2, and Ce13, respectively; all other test solutions were obtained by serial dilution of these stock solutions. Procedure. Potentiometric measurements for the construction of the calibration curves were carried out in thc conventional manner. I n every case, the electrode system was allowed to attain equilibrium as indicated by a stable (AO.1 mv.) cell e.m.f. Repetitive measurements, carried out for each data point, were reproducible to better than 1 mv. when performed on the same day. Potentiometric titrations were carried out using the incremental addition of standardized silver(1) solutions (prepared from XgC104) to sample solutions containing known amounts of KI.
-100
1
ii u 0
g
1.0
KI
EoIz A CeI,
2.0 3.0 4.0
5.0
6.0 7.0 8.0
- l o g [I-]
Figure 3. Test of cation effect on electrode response. Potentials measured vs. saturated calomel electrode
The titration of silver(1) with iodide is undesirable because it exposes the indicator electrode to a possible loss of halide by reaction with excess silver(1). Erratic e.m.f. values, depending on exposure time and stirring efficiency, may be encountered under such conditions. Electrode response curves were obtained after rapid dilution or concentration of the test sample with water or electrolyte solution, respectively. Because of mixing and iiistrumental limitations, no meaningful response data could be obtained a t timeb less than 1 to 2 seconds after mixing. In all cases, the resulting e.m.f. us. time curves were smooth and reproducible after the initial mixing period. RESULTS AND DISCUSSION
Ip
2p
50
40
-log
5D 60 7.0 E.(
[I-]
Figure 2. Calibration curve for membrane electrode. Concentrations are molar concentrations of potassium iodide. Potentials measured vs. saturated calomel electrode
Figure 2 shows a typical calibration curve constructed from potentiometric data obtained with one of the membrane electrodes in KI solutions. The remarkable sensitivity and dynamic range of the electrode is evident from the fact that the calibration curve is linear over the KI concentration range of 10-l to nearly 10-731. The erratic results obtained a t concentrations below 5 X l O - 7 M are not surprising in view of the possible presence of impurities and the difficulty of preparing standard solutions a t the submicromolar concentration level. The slope of the curve in Figure 2 is 55.3 mv. per concentration decade; while all of the iodide electrodes tested gave such linear calibration curves, the numerical values of the resulting slopes ranged from a low of about 53 mv. per
+600
t 500
+400 >' E
ti
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t 300
i
t 200
+
J
i-
*
t 100
++,-+' 4 /-Y+
0
1
I
I
I
I
I
I
I
2
3
4
5
6
7
8
ml. 0 . 0 1 1 Ag+
Figure 4. Potentiometric titration using membrane indicator electrode. Potentials measured vs. saturated calomel electrode
decade to a high of 59.1 mv. (the Nernstian value). The reproducibility of individual e.m.f. measurements is only of the order of *0.5 mv., however, so that calibration curves are best constructed by averaging sets of repetitive measurements. Electrode drift is not a significant contribution to the overall error during the time required for construction of a calibration curve, but necessitates recalibration of the electrode system every one or two days. Potentiometric equilibrium was always attained (except for samples
