(6) M. M. Bursey. T. A. Elwood, M. K. Hoffman, T. A. Lehman, and J. M. Tesarek, Anal. Chem., 42, 1370 (1970). (7) E. L. Eliel, "Stereochemistry of Carbon Compounds", McGraw-Hill Book Company, New York, NY, 1962, p 219. (8) E. L. Eliel, "Stereochemistry of Carbon Compounds", McGraw-Hill Book Company, New York NY, 1962, p 222. (9) H. R. Nace and R. H.Nealey, J. Am. Chem. SOC., 88, 65 (1966). (10) R. L. Stern, J. S. Zannucci, and 8. L. Karger, Chem. Commun., 613 (1967). (1 1) R. W. J. LeFevre, Adv. fhys. Org. Chem., 3, 1 (1965). (12) G. Gioumousis and D. P. Stevenson, J. Cbem. fbys., 20, 294 (1958). (13) T. F. Moran and W. Hamill, J. Chem. Phys., 39, 1413 (1963). (14) T. Su and M. T. Bowers, J. Chem. fhys., 58, 3027 (1973). (15) T. Su and M. T. Bowers, Int. J. Mass Spectrom. /on Phys., 12, 347 (1973). (16) R. C. Dunbar, M. M. Bursey, and D. A. Chatfield, Int. J. Mass Spectrom. Ion Phys., 13, 195 (1974).
M. L. Gross and J. Norbeck. J. Chem. fhys., 54, 3641 (1971). M. K. Hoffman and M. M. Bursey, Can. J. Chem., 49,3995 (1971). S. A. Benezra and M. M. Bursey, J. Am. Chem. SOC., 94, 1024 (1972). J. D. Henion, M. C. Sammons, C. E. Parker, and M. M. Bursey, Tetrahedron Lett., 4925 (1973).
RECEIVEDfor review December 9, 1974. Accepted March 10, 1975. We are grateful for support from the National Institute of General Medical Sciences (GM 15,994). The instrument was purchased from funds donated by Hercules, Inc., the Shell Companies Foundation, the North Carolina Board of Science and Technology, and the National Science Foundation (GU 2059).
Application of Ion Exchange Membranes to Sampling and Enrichment: Interference of Metal Ion Binding Groups Walter J. Blaedel and Richard A. Niemann Department of Chemistry, University of Wisconsin, Madison, WI 53706
The identification or quantitation of cations present a t trace concentration levels in natural water systems has often involved an enrichment step prior to the analysis ( I ) . Ion exchangers are particularly well-suited for such enrichment, and many applications have been described in which the enrichment occurs into the ion exchanger phase. Materials have been particulate or membranous, and chelating resins also have been used (2-9). In principle, ion exchange membranes may be used in a different mode with particular advantage as sampling-enrichment devices, by concentrating the trace metal ions from a large volume of a dilute water sample (the donor solution) through the ion exchange membrane into a small volume of a concentrated solution (the acceptor solution). Measurement of the enriched cations in the acceptor solution would be simpler than in the original donor solution. The particular advantage of such a system for sampling and enrichment purposes is that the trace cation distribution depends only on the charge types of the ions, and is independent of the other physical and chemical properties of the ions or of the membrane. The theoretical basis of the distribution is well established, and it has been demonstrated experimentally (20-13). Exploratory attempts to utilize this approach for sampling and enrichment of Cu(I1) a t micromolar concentration levels required very long equilibration times, and led to the hypothesis that the membrane contained impurity groups capable of binding the metal ion very strongly by mechanisms other than the ion exchange mechanism. The work described in this note was undertaken to demonstrate unequivocally the occurrence of such binding, to establish the levels at which it operates for some of the commercially available membranes, and to seek infrared spectrometric evidence supporting the presence of impurity groups in the membranes.
EXPERIMENTAL Reagents. All chemicals used were reagent grade. All solutions were prepared from triply distilled water: once-distilled t a p water
was redistilled from alkaline permanganate, collected, made about 0.001M in HzS04 and distilled again. For making dilute solutions, an aqueous stock solution of 0.01M Cu(N03)2 was prepared and standardized by titration with EDTA. Solutions of lower concentration were prepared by appropriate dilution of aliquots of the stock solution. Radiotracer 64Cu was prepared in the University of Wisconsin nuclear reactor facility by thermal neutron irradiation of a weighed amount (about 30 mg) of pure metallic copper. After irradiation, the copper was dissolved in a few milliliters of 8M "03 and evaporated to near dryness over a low flame. The residue was evaporated with about 20-ml portions of water twice more to near dryness, and finally transferred with 0.10M NaCl to a 100-ml volumetric flask and made u p to volume with 0.10M NaC1. An aliquot of this stock solution was used in the first extraction of the successive extraction experiments. Three types of sulfonated cation exchange membranes were studied: AMF (2-103 (American Machine and Foundry Co., Stamford, CT, sulfonated styrene on a polyethylene backbone, 0.007 inch thick, 1.3 mequiv per dry gram of H-form); Nafion XR-170 (Plastics Dept., Du Pont and Co., Wilmington, DE, sulfonated fluorocarbon polymer, 0.0035 inch thick, 0.83 mequiv per dry gram of H-form); Permion PlOlO (RAI Research Corp., Hauppauge, Long Island, NY, T F E Teflon-sulfonated styrene copolymer, 0.0015 inch thick, 2.00 mequiv per dry gram of H-form). Before use, all membranes were conditioned and converted to the Na-form by the following steps: 1) wash for 0.5 hour in 95% ethanol; 2) three cycles of 5-minute rinses in 1M HCI, water, and 1M NaOH; 3) three half-hour washes in portions of 1M NaC1; 4) storage in 0.10M NaCl until needed. Copper(I1) Distribution Measurements. A weighed rectangular piece of membrane (about 0.1 mequiv total ion exchange capacity) was contacted with mechanical stirring for about 8 hours with a measured weight (between 150 and 500 g) of 0.10M NaCl solution that contained a known concentration of copper (about 15 ppm, or about 0.2mM), and that also contained r a d i ~ - ~ ~(activiCu ty greater than lo4 cpm/g solution). A significant fraction (5 to 10%) of the copper was taken up by the membrane, and the distribution ratio and copper content of the membrane were determined from the counting rate of the solution before and after extraction. The membrane was removed from solution, blotted dry with a tissue, and counted for 64Cu activity. T o ensure a reproducible counting geometry, the membrane, wrapped around a Plexiglas rod having a slit in one end, was inserted into the bottom of the counting tube where it was uncoiled from the rod and carefully positioned ANALYTICALCHEMISTRY, VOL. 47, NO. 8, JULY 1975
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3f
u
AMF C-103
1.000
OUFUNT N A F I W XR-170
I 1.0 100 EQUILIBRIUM COPPER CONTENT OF MEMBRANE, m e g 7.
Flgure 1. Distribution ratio of Cu(ll) between various ion exchange membranes and 0.10M NaCl ('Ordinate values for AMF C-103 are ten times larger than indicated)
around the inner circumference of the tube. The rod was then withdrawn from the tube. A second extraction was made under conditions identical to the first, except that copper-free 0.10M NaCl was used as the extractant. Counting of the membrane before and after extraction permitted calculation of the distribution ratio and copper content of the membrane. Successive extractions were performed until the activity became too low to count accurately. Because of the duration of the experiments, corrections were applied for decay. In fact, for the last few extractions, rather long counts were usually taken which required a correction for decay during the counting interval (14). Coincidence corrections were applied to counting rates over lo5 cpm. Counting was done in a scintillation counter (Baird Atomic, Chicago, IL= Model 530 spectrometer; Model 810C well-type scintillation detector (1.75-inch X 2-inch NaI(T1) crystal); and Model 620 printer). The pulse height analyzer was adjusted to accept only those counts originating from positron annihilation (511 keV yray). The radiotracer technique was used on the Nafion and Permion membranes. The AMF membrane had a much higher level of impurity groups, and atomic absorption was used in measurement instead of radioactivity. Infrared Spectra. Transmission infrared spectra were taken on separate pieces of membrane in the sodium form, using a PerkinElmer Model 421 dual grating double beam spectrophotometer. Spectra obtained with the hydrogen forms of the ion exchange membranes were not interpretable.
RESULTS AND DISCUSSION Distribution Ratio Measurements. Figure 1 is a plot of the distribution ratio of copper vs. the copper content of the membrane, expressed as mequiv percent. A log-log plot is used because of the wide range covered by the multiple extractions. If only the ion exchange mechanism operated to bind copper, its distribution ratio would approach a constant value a t low copper contents, since the sodium ion molarity in all extractions is 0.10M. However, a t a copper content where a good fraction of the bound copper is held tightly by trapping groups in the membrane, the distribution ratio would rise to a high value. Figure 1 indicates that the trapping group contents of the AMF (2-103, Nafion XR-170, and RAI PlOlO cation exchange membranes fell around 5, 0.01, and 0.05 mequiv %, respectively. Infrared Spectra. In general, the major functional groups in the membrane matrix (including atmospheric and membrane-bound water) caused regions of absorption in which impurity group absorption could not be observed. As a consequence, impurity ;group detection by infrared spectrometry was by no means complete, but only indicative. Also, n o attempt was made to quantitate the relative importance of the detected impurity groups in actually binding Cu(I1). The AMF membrane exhibited a broad band of moderate intensity in the 1725-1700 cm-l region (carbonyl) and weak bands at 1565 cm-l (carboxylate anion stretch) and 920 cm-I (vinyl C-H bend). T o test for the presence of carboxylic acid groups, spectra were obtained on a series of 1456
ANALYTICAL CHEMISTRY, VOL. 47, NO. 8, JULY 1975
membranes in which the H+ content was increased gradually. With increasing H+ content, the 1700-1725 cm-l band shape became progressively more definitive and intense while, a t the same time, the intensity of the 1565 cm-l peak decreased (less carboxyl content in the anion form). Reacting the membrane for one day in a bromine-sodium bromide aqueous solution resulted in a decrease in the 920 cm-l peak intensity as well as substantial increases in both the 1700-1725 cm-l and 1565 cm-I intensities. This was interpreted to mean that bromine treatment not only saturates impurity olefinic bonds through bromination but also generates additional carbonyl (ketonic and aldehydic) and carboxylate groups. The Nafion membrane showed moderately intense peaks a t 2990, 2920, 1462, 1415, and 888 cm-' (olefinic hydrocarbon). Substantial reduction in these peak intensities was achieved by reacting the membrane for one day in an aqueous bromine-sodium bromide solution. Even greater reductions were obtained with a five-day reaction period. No definite impurity peaks were detected in the RAI membrane, possibly due to the very short path length through the thin membrane. Other Evidence for Impurity Groups. The AMF membrane exhibited a color change (light brown to a darker brown) on transition from the acid to the salt (base) form, indicating the presence of chromophoric (conjugated) systems. Rieman and Walton (15) have stated that rarely is an ion exchanger truly monofunctional. Furthermore, the chemical literature contains evidence for the presence of carboxyl groups in sulfonic acid resins (16-18) and for sulfone groups in some sulfonic acid membrane preparations (19). An attempt to use an ion exchanger to determine the hydrolysis constants of indium a t submicromolar concentration levels was unsuccessful due to the presence of impurity groups that strongly bind indium (20).
CONCLUSIONS I t has been demonstrated that ion exchange membranes of the sulfonic acid type contain impurity groups, some of which have the capacity to bind copper ion strongly. The presence of these impurity groups interferes with normal ion exchange behavior a t low concentration levels of copper ion, and may preclude use of the membrane in some procedures for trace analysis. ACKNOWLEDGMENT The assistance of Robert Schmelzer and Robert Lang in machining operations is highly appreciated. Gifts of ion exchange membrane materials from E. I. du Pont de Nemours and Co., and from RAI Research Corp. are also greatly appreciated.
LITERATURE CITED (1)
J. B. Andelman and S. C. Caruso, Water Water Polut. Handb., 2, 483
(1971). (2) S. L. Law, Science, 174, 285 (1971). (3) J. F. Dinaman. Jr.. K. M. Gloss, E. A. Milano. and S. Siaaia. Anal. Chem., 46, 774 Ti974) (4) D E Becknell, R H Marsh, and W Allie. Jr , Anal. Chem., 43, 1230 11971) -. I (5) R. G. Smith, Jr., Anal. Chem., 46, 607 (1974). (6) 0. Samuelson, "Ion Exchange Separations in Analytical Chemistry", Wiley, New York, 1963, pp 293-298. (7) W. Riernan 111 and H. F. Walton. "Ion Exchange in Analytical Chemistry", Pergamon Press, New York, 1970. (8) U. Eisner and H. B. Mark, Jr., Talanta, 16, 27 (1969). (9) C. H. Lochmuller, J . W. Galbraith, and R. L. Walter, Anal. Chem., 46, 440 (1974). (10) W. J. Blaedel and T. J. Haupert. Anal. Chem., 38, 1305 (1966). (11) R. M. Wallace, lnd. Eng. Chem., ProcessDes. Develop., 6, 143 (1967). (12) W. J. Blaedel and T. R. Kissel, Anal. Chem., 44, 2109 (1972). (13) K . N . Pearce and L. K. Creamer, Anal. Chem. 46, 457 (1974).
--
I
Aziz and S.J . Lyle, J. lnorg. Nucl. Chem., 31, 2431 (1969).
(14) D.
(20) A.
(15) Ref.'7, p. 51. (16) I. M. Abrams, lnd. Eng. Chem., 48, 1469 (1956). (17) J. Lindeman, zesz, Nauk. Politech, Lodz., wlok 1969, NO. 204, 3. (Chem. Absfr., 71, 39693h (1969)). (18) G. M. Armitage and S. J . Lyle, Talanta. 20, 315 (1973). (19) C. W. Plummer, J. Enos. A. 8 . LaConti, and J . R. Boyack, Office of Saline Wafer, Res. Develop. Rep. No. 481 (1969).
RECEIVEDfor review November 18, 1974. Accepted March 24, 1975. This research was supported in part by an Office of Water Resources Research Grant, No. A-053-WIS. Addiin the form Of a Du Summer Research Assistantship (R.A.N., 1972) is gratefully acknowledged.
deSoete, R. Gijbels. and J. Hoste, "Neutron Activation Analysis", Wilev-Interscience, London, 1972, ,DD. 521-523.
Catalyst Electrode Specific for Peroxide Stuart J. Updike and Mark C. Shults Department of Medicine, University of Wisconsin, Madison, Wl 53706
Judy K. Kosovich Department of Chemical Engineering, University of Wisconsin, Madison, Wl 53706
Isabel Treichel and Paul M. Treichel Department of Chemistry, University of Wisconsin, Madison, Wl53706
Polarographic measurement of 0 2 and H202 is the basis for several enzyme-coupled electrodes (1-3). Hydrogen peroxide can be polarographically assayed by anodic oxidation (2).This method is very sensitive, but lacks specificity for peroxide in the presence of other redox-active substances that undergo anodic oxidation, and which are present in biological solutions, for example, ascorbic acid. However, if H202 is first catalytically decomposed to 0 2 and H20, then the 0 2 can be assayed polarographically using an oxygen electrode. Furthermore, specificity for just oxygen can be obtained by placing a hydrophobic membrane over the 0 2 electrode, as originally described by Clark (2).Oxygen is the only electroactive substance in biological solutions that can cross a hydrophobic membrane and generate an electrode current. Herein we report an inorganic catalyst electrode. This peroxide sensing device is made by covering an oxygen electrode with a membrane that catalyzes the breakdown of hydrogen peroxide to oxygen. This electrode is similar in principle to an enzyme electrode (1, 3), but has the added stability gained by using an inorganic catalyst, rather than an enzyme such as catalase.
EXPERIMENTAL T h e oxygen sensor used to build the catalyst electrode was a Clark type of microcathode oxygen electrode (Radiometer, Copenhagen, Denmark, Type No. E-5046). The current output of this type of electrode is a linear function of oxygen tension ( 4 ) . A membrane made of regenerated cellulose (Neflex hemodialysis membrane, 20-micron wet thickness, Union Carbide, Chicago, IL) is modified so as to carry out transmembrane catalysis of H202. This membrane is press-fitted over the oxygen electrode with an O-ring. Catalyst Impregnated Membrane. Noble and transition metals form insoluble oxides and sulfides, many of which are wellknown catalysts for breaking down H202 to 0 2 and H20 ( 5 ) .These catalysts can be precipitated within the polymer matrix of the cellulosic dialysis membrane by contacting one side of the membrane with 0.3M chloride salt solution of the transition metal and contacting the other side of the membrane simultaneously with 0.1M sodium hydroxide or 0.1M sodium sulfide. A contact time of 1 minute was ordinarily used.
Initial tests for catalytic activity were made on membranes in which the following metal compounds were deposited: Fe203, FeS, NiO, NiS, Ag;?O, AgzS, CuO, CuS, Zn(OH)2, ZnS, Cr203, Cr2S3, Cd(OH)2, CdS, La(OH)3, CeO2, MnO2, Ru203, Ru&, C0203, COS. All but Zn(OH)p, ZnS, Cd(OH)2, CdS, La(OH)Z, and CeOz showed some degree of catalytic activity; these results parallel the results of earlier (non-membrane) studies ( 5 ) . (The MnO2-containing membrane was prepared by contacting the membrane with 0.3M KMn04 and 0.1M NaI simultaneously for 1 minute on opposite sides of the membrane. Reduction of permanganate by iodide produced a precipitation of MnO2 in the membrane. The CozO3-containing membrane was formed by first contacting the membrane simultaneously with 0.3M CoC12 and 0.1M NaOH to precipitate blue Co(OH)2. Upon contact with H202, the cobalt hydroxide was oxidized t o greenish-black hydrated CozO3.) Of the above listed 21 different metal compounds t h a t were deposited in cellulosic membrane, only the compounds of manganese, cobalt, and ruthenium showed high catalytic activity. Catalytic activity was measured semiquantitatively by comparing the vigor with which 0 2 bubbles are produced on the membrane surface in a 0.2% solution of H202. Catalytic activity was measured quantitatively by a volumetric method detailed in a previous report (6). Only compounds of these three metals deposited in membranes were further studied for stability and washout of the catalyst.
RESULTS AND DISCUSSION Catalyst Stability with Respect to pH and Chelation. Gamma-ray emitting isotopes of manganese, cobalt, and ruthenium were available to us through the neutron activation services of the University of Wisconsin Department of Nuclear Engineering. Catalyst washout as a function of pH was determined by observing any washout from the membrane as a function of pH in the range of 2 to 10 in the presence of H202 kept a t a concentration of 0.2 to 0.5%. A summary of data on pH stability for the oxides of manganese, cobalt, and ruthenium is given in Table I. The oxide and sulfide of a given metal behaved identically. The oxide and sulfide of cobalt and the oxide of manganese were excellent catalysts in alkaline solution. Over a 24-hour period there was less than 3% 56Mn washout above pH 9.0, and less than 3% 6oCo washout above pH 8.0. However, both of these metals washed out of the membrane completely within two hours at acid pH. ANALYTICALCHEMISTRY, VOL. 47, NO. 8, JULY 1975
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