@CH3CON133ci&
There are many possible types of applications inherent in an electrochemical pH-stat. Through the use of immobilized enzymes, the rapid analysis of nonelectroactive materials may be performed with coulometric precision and accuracy. Work is now under way to apply the instrument to purity determinations of nonelectroactive materials and to the development of a flow system for use in clinical analysis.
COzH
PENICILLIN G
Hb PENICILLOICACID The penicilloic acid is the substance coulometrically titrated in the pH-stat. Table VI illustrates the measurement of a series of penicillin G samples. This compound is hygroscopicand heat sensitive making it difficult to weigh accurately. The precision obtained in Table VI indicates that the pH-stat is capable of coulometrically analyzing small amounts of penicillin G and the method should be applicable to a great many additional materials. As indicated above, the system cannot be applied to a first-order chemical reaction whose half-life is less than 15 s. A convenient test system is the well characterized hydrolysis of tert-butyl chloride in water at pH 7.0: CHB
I CH3-C-C1 I
+
H20
-
CH3
I
CH,-C-OH
I
+
HCl (20)
CH3
CH3
A typical coulombs-time curve for this reaction is shown in Figure 10 along with a first-order rate plot. An excellent fit was obtained (correlation coefficient = 0.99; Student’s t = 100) and a rate constant of 0.0408 s-1 obtained from the average slope of 5 runs. The between run relative standard deviation was 2%.The above value is in agreement with the estimates of Herbrandson (17) which ranged from 2.0 to 3.7 X s-l. This reaction half-life is almost at the upper limit of the system and indicates that the system can be used to measure rates in this region with acceptable accuracy and precision.
ACKNOWLEDGMENT The receipt of a sample of Nafion 125 ion exchange membrane from E. I. du Pont de Nemours and Company is gratefully acknowledged. LITERATURE CITED (1) C. F. Jacobsen, J. Leonis, K. Linderstrom-Lang, and M. Ottesen, “Methods of Biochemical Analysis”, D. Glick, Ed., lnterscience Publishers, New York, N.Y., 1957, Vol. IV, p 171. (2) D.W. Elnsel, Jr., H. J. Trurnlt, S. D. Silver, and E. C. Stelner, Anal. Chem., 28,408 (1956). (3) R. E. Karcher and H. L. Pardue, Clin. Chem. ( Winston-Salem, N.C.),17, 214(1971). (4) H. V. Malmstadt and E. H. Piepmeier, Anal. Chem., 37,34 (1965). (5)L. Meites and H. C. Thomas, “Advanced Analytical Chemistry”, McGraw-Hill Book Co., New York, N.Y., 1958, p 73. (6) 0. Johansson, Talanta, 12, 111 (1965). (7) M. J. D. Brandand G. A. Rechnitz, Anal. Chem.. 41, 1185 (1969). (8)J. A. Page and J. J. Lingane, Anal. Chim. Acta, 16, 175 (1957). (9) J. E. Harrar and R. J. Sherry, Anal. Chem., 47, 601 (1975). (IO) A. Wikby and B. Karlberg, Electrochim. Acta, IS, 323 (1974). (11) B. Karlberg, J. Electroanal. Chem., 49, 1 (1974). (12) D.R. Thatcher, “Methods in Enzymology”, J. H. Hush, Ed., Academic Press, New York, N.Y., 1974, Vol. 43, pp 640-687. (13) R. Labia, J. Andrillon, and F. LeGoffic, Biochem. Biophys, Acta, 384,242 (1975). (14) J. P. Houand J. W. Poole, J. Pharm. Sci., 61, 1594(1972). (15) J. J. Martaugh and G. B. Levy, J. Am. Chem. Soc., 67, 1042 (1945). (16) W. S. Wise and G. H. Twig, Ana/yst(London), 75,106 (1950). (17) H. F. Herbrandson, J. Chem. Educ., 48,708 (1971).
RECEIVEDfor review March 22, 1976. Accepted August 4, 1976. This work was supported by Grant 17913from the National Institutes of Health. Presented in part at the Pittsburgh Conference on Applied Spectroscopy and Analytical Chemistry, Cleveland, Ohio, March 1-5, 1976.
Direct Potentiometric Measurement of Sulfite Ion with Mercuric Sulfide/Mercurous Chloride Membrane Electrode Paul K. C. Tseng and W. F. Gutknecht” Department of Chemistry, Duke University, Durham, N.C. 27706
An electrode having an ion-sensing membrane composed of a polycrystallinemixture of HgS and Hg2CI2has been used for the direct potentiometric measurement of the sulfite Ion. It Is proposed that the response mechanlsm involves the reaction Hg2CI2 2S032- -,Hg Hg(S03)22- 2CI-, and Its effect upon the equllibrium Hg2CI26 Hg22+ 2CI-. The llmlt of detection for this measurement is below 1 part-per-million.
+
+
+ +
A number of methods are available for the determination of the sulfite ion (or sulfur dioxide) (1-13). Among these are a method based on the redox reaction between sulfite and iodine ( 1 )as well as several spectrophotometric methods (6, 8, 9, 11). Two potentiometric methods are available for the analysis of the sulfite ion. The first involves the use of the double-membrane electrode; a gas-permeable membrane 1996
separates the sample solution from an internal buffer solution (sodium bisulfite) in contact with a pH-sensing glass electrode (12). Sulfur dioxide in equilibrium with the sulfite in the sample solution diffuses through the membrane, interacts with the sodium bisulfite of the internal solution so as to change its pH, and this pH change is sensed by the glass electrode. Electrodes of this type are available from several commercial sources. An iodide-sensitive electrode has been used indirectly for sulfite analysis (13). In this method, sulfite is first oxidized to sulfate with iodine, after which the iodide formed is measured using the silver iodide-based ion-selective electrode. We are reporting herewith on the application-ofthe HgS/HgZC12 membrane electrode (as first developed by Lechner and Sekerka (14,15) for the analysis of chloride) to the analysis of the sulfite ion. As will be discussed in detail below, the mercurous chloride in this membrane is the active component whereas the mercuric sulfide serves as a binder of
ANALYTICAL CHEMISTRY, VOL. 48, NO. 13, NOVEMBER 1976
300
dg
-
200.
100.
Y
210
I
2
I
4
I
6
I
8
I
10
PH
Flgure 2. Response of the HgS/Hg2CI2 membrane electrode to M sulfite ion as a function of test solution pH
Figure 1. Response of the HgS/Hg2CI2membrane electrode to sulfite ion. Test solution pH 5-6
low solubility and high conductivity (14). In general terms, the sulfite ion reacts with the mercurous chloride in the membrane to form a soluble complex. This response mechanism is not unlike that of the AgZS/AgI membrane response to cyanide ion (16-18).
Table I. Comparison of the Response Potentials of Sulfite and Chloride at Different Concentrations Concentration,M Test solution
Sulfite ion Chloride ion
10-3
131 mV 131 mV
180 mV
181 mV
10-4 230 mV 231 mV
10-5
282 mV 281 mV
THEORY
It is well known that several moderately insoluble mercury(1) compounds can react with certain anions to form stable complexes of mercury(I1)and elemental mercury (19).Sulfite ion is one such species which causes the disproportionation of mercury(1) (11,20). That is, Hg2C12
+ 2S03'-
+ Hg
+ Hg(S03)Z2- + 2C1-
(1)
This reaction has been qualitatively described as going essentially to completion ( I I ) , though apparently no quantitative expression for this reaction is available. Thus such an expression has been calculated. Now one may view the reaction as a three-step process; that is, Hg 2a2
+
Hgz2+ 2C1-
Hgz2++ Hgo
+
Hg2+ 2So&
+ Hg2+
K,, = 10-17.88 (21) Kdis = 10-2.22(22)
Hg(S03)z2-
(2) (3)
Kf = lOZ4.O7 (23) (4)
[The Kdis for Reaction 3 was calculated from the standard potentials for the reactions Hgz2+ * 2Hg2+ + 2e- (Eo = -0.920 V) and Hgzz+ 2e2Hg0 (Eo = 0.789 V) (22).] Utilizing these three expressions, an equilibrium expression for Reaction 1is calculated, i.e.,
+
Ke, =
+
(Hg(SO3)z2-)(Cl1-P (S032-)2
= (Ks,)(Kdis)(Kf) = 103.97
(5)
With this large driving force for the reaction, each sulfite ion will displace one chloride ion at the surface of the electrode, and as the electrode is aparently responding to the chloride ion via Reaction 2, one will obtain identical responses to both chloride and sulfite ions (see below). EXPERIMENTAL The HgS/HgZClz membrane electrodes used in this work were prepared in our laboratory; their fabrication is described in Reference 14.Mercuric sulfide (black form) is prepared by adding a slight molar excess of either NazS or thioacetamide solution to a solution of mercuric nitrate. The mercuric sulfide so prepared is well mixed oneto-one by weight with mercurous chloride (Baker Chemical Co.) using a mortar and pestle. Then 1-g portions of this mixture are compressed at 27 000 psi in a KBr press to form membranes 13 mm in diameter and about 1mm thick. The membranes are attached to a glass tube with epoxy and internal contact is made with a small quantity of mercury (14). The reference electrode used was a saturated calomel electrode. Connection between the reference electrode and test so-
lution is made with a fiber-tipped secondary junction filled with 1M KNO3. These laboratory-made electrodes show a linear response for E,,,, vs. log (C1-) wherein the response is 50-51 mV per decade of chloride ion concentration. Though numerous attempts were made to do so, it was not possible to reproduce the 58-59 mV per decade response obtained by Lechner and Sekerka (14). Water used in this work was purified by deionization followed by distillation. All chemicals used were reagent grade. The sodium sulfite test solutions were made up in water containing about 5% glycerol to prevent oxidation ( 1 1 , 2 4 ) . All cell emf measurements were made with a Beckman SS-3 pH meter. Solutions were stirred during each measurement and the solution temperatures were ambient at 20-21 "C. The final potential of each test solution was determined when the measured potential changed less than 0.5 mV over a 3-min period. Replicate measurements were generally reproducible to within 10.5 mV. Final cell voltages were reached with chloride solutions in about 15 s; sulfite solutions required about 30 s for a final cell voltage to be reached. The electrode membranes were not polished before each measurement though they were polished after initial mounting on the glass tube, and they were soaked several hours in water before they were used.
RESULTS AND DISCUSSION
First to be established was the linear response range. Figure 1shows this range to be 0.05 to 10-5 M with a response slope of 50 mV per decade. Above 0.05 M, the response is no longer linear. This is due to a change in ionic strength of the test solutions with concentration (see below) and also the formation of a second mercury(I1)-sulfite ion complex, Hg( (20). A t this high concentration of sulfite ion, the chloride ion is displaced on less than a one-to-one basis by sulfite ion and therefore the response potential is higher than expected. Support for Reaction 1playing a major role in the electrode response mechanism is given in Table I. Here the electrode potentials for a series of chloride ion and sulfite ion test solutions are compared. Their near-equivalencies support the supposition that the response mechanism is indeed the displacement of chloride ion by sulfite ion with a subsequent alteration in the equilibrium between Hg2C12, Hgzz+and C1-. This response is, of course, analogous to the response of the Ag2S/AgI membrane electrode to cyanide ion (25-27). Now another possible response reaction to be considered is Hg2Clz + SO2
+ 2H20* 2Hg + S042- + 4H+ + 2C1-
(6)
This reaction produces, however, two chloride ions from each single SO2 entity and thus does not explain the values found in Table I. More importantly, the equilibrium expression for
ANALYTICAL CHEMISTRY, VOL. 48, NO. 13, NOVEMBER 1976
lQQ7
Table 11. Selectivity Coefficients for Several Ions Interfering with Sulfite Measurement Using the HgS/ HgZC12 Membrane Electrode Interfering ion C1-
Br-
I-
K~~~ 1 102.1 105.0 a Determined at pH 5-6.
SCN-
NO3-“
C1041-a
100.8
,-10-3
-10-3
this reaction was calculated as for Reaction 1and is found to be only 10-7.48.This value indicates clearly that Reaction 6 has no significant role in the production of chloride ion when compared to Reaction 1and thus does not enter into the response mechanism. The pH value of the sulfite test solution is an important parameter to be considered. Figure 2 shows the electrode response to a M test sulfite solution as a function of solution pH. A useful range appears to be 3.3 to 8.5. Now much of the test ion is in the bisulfite form throughout this pH range. Nevertheless the electrode is responding to the sulfite ion in that the hydrogen ion is a poor competitor with the mercury(I1) ion for the sulfite ligand; that is,
+
Hg2+ 2SOs2-
* Hg(S03)z2-
K = 1024.07(23)
(8)
When the pH is greater than 8.5, hydroxide ion apparently interferes through the formation of Hg2(0H)2 (Ksp= (21 1.
When the pH value of the test solution drops below 3.3, sulfur dioxide becomes the major solution component present as the existence of sulfurous acid (H2SO3) is still unknown (24). This fact explains the increase in the response potential at the lower pH values. Attempts were made a t controlling the ionic strength of the test solutions using KN03, NaN03, and NaC104. All three interfered with the sulfite measurements. These substances most likely interfere through the formation of the hydroxy mercurous salts HO.Hg2N03 (29,30) and HO.Hg2C104 (31). This rationale is supported by the fact that the degree of interference from these species decreases with decreasing test solution pH. Because of these interferences, our measurements were made without ionic strength control; Lechner and Sekerka did likewise (14). In addition to nitrate and perchlorate, other interfering species include bromide, chloride, cyanide, and iodide ions, and also some thiols (14). The selectivity coefficients for these ions with respect to sulfite ion were determined using a standard approach (32), and are shown in Table 11. One will note that the halide interferences are not unlike those found
1998
with the AgzS-based membrane electrodes. The interference due to nitrate and perchlorate is not large and could be reduced by lowering the test solution pH. In conclusion then, the HgS/HgzC12 membrane electrode has been successfully used for the potentiometric measurement of the sulfite ion, albeit the number of potential interferences poses limitations on its application. More importantly a response mechanism has been proposed. This proposed response mechanism and supporting data about an electrode other than a AgzS-based membrane electrode are valuable in that they add to the pool of general knowledge regarding the behavior of solid-state, ion-selective electrodes. LITERATURE CITED (1) “Scott’s Standard Methods of Chemical Analysis”. 5th ed., N. H. Furman, Ed., Van Nostrand, Princeton, N.J., 1939, p 926. (2) A. Steigmann, J. Soc. Chem. lnd., London, 61, 18 (1942). (3) W. M. Grant, Anal. Chem., 19, 245 (1947). (4) S.Atkin, Anal. Chem., 22, 947 (1950). (5) P. F. Uron and W. E. Boggs, Anal. Chem., 23 (1951). (6) W. L. Hinze and R. E. Humphrey, Anal. Chem., 45, 385 (1973). (7) W. L. Hinzeand R. E. Humphrey, Anal. Chem., 44, 1511 (1972). (8)T. Okutani and S.Utsumi, Bull. Chem. SOC., Jpn., 40, 1386 (1967). (9) T. Okutani, S. Ito, and S. Utsumi, Nippon Kagaku Zasshl, 88, 1296 (1967). (IO) R. J. Pfeifer, B. Y. Cho, and 0. L. Utt, /SA Trans., 9, 9 (1970). (11) W.L. Hinze, D. J. Kippenberger, and R. E. Humphrey, Mlcrochem. J., 20, 43 (1975). (12) J. W. Ross, Jr., J. H. Riseman, and J. A. Krueger, J. Pure Appl. Chem., 36, 373 (1973). (13) S. Ikeda, J. Hirata, and H. Satake, Nlppon K8gaku Zasshl, 94, 1473 (1973). (14) J. F. Lechner and i. Sekerka, J. Electroanal. Chem. lnterfacialElectrochem., 57,317 (1974). (15) I. Sekerka and J. F. Lechner, J. Electroanal. Chem., 69, 339 (1976). (16) K. Toth and E. Pungor, Anal. Chim. Acta, 51, 221 (1970). (17) B. Fleet and H. Von Storp, Anal. Lett., 4, 425 (1971). (18)M. S. Frant, J. W. Ross, Jr., and J. H. Riseman, Anal. Chem., 44, 2227 (1972). (19) F. A. Cotton, and G. Wilkinson, “Advanced inorganic Chemistry”, 22d ed., lnterscience Publishers, New York, 1967, p 613. (20) B. J. Ayiett, “Comprehensive Inorganic Chemistry”, Vol. 3, Compendium Publishers, New York, 1973, p 316. (21) A. Ringbom, “Complexation In Analytical Chemistry”, Wiiey, New York, 1963, p 343. (22) E. J. King, ”Qualitative Analysis and Electrolytic Solutions”, Harcourt, Brace &World, New York, 1964, p 610. (23) Chemical Society Special Publication No. 17, “Stability Constants of Metal Ion Complexes,” London, 1964. (24) P. J. Durrant and B. Durrant, “Introduction To Advanced inorganic Chemistry”, 2d ed., Wiey, New York, 1970, p. 849. (25) B. Fleet and H. Von Storp, Anal. Chem., 43, 1575 (1971). (26) D. H. Evans, Anal. Chem., 44, 875 (1972). (27) M. Mascini, Anal. Chem., 45, 614 (1973). (28) Reference 22, p 605. (29) P. C. Ray, J. Chem. Soc. Trans., 87, 171 (1905). (30) “Gmelins Handbuch Der Anorganischem Chemle”, Teil B, Lieferung I, System Number 34, pp 95-114, 1965. (31) Reference 22, p 610. (32) Paul K. C. Tseng and W. F. Gutknecht, Anal. Lett., in press.
RECEIVEDfor review April 1,1976.Accepted August 9,1976. The authors wish to express their sincere appreciation to the Duke University Research Council for its financial support of this work.
ANALYTICAL CHEMISTRY, VOL. 48, NO. 13, NOVEMBER 1976
