data were available to set limits of error on the constants, but a reasonable estimate is less than 1 0 . 5 logarithmic unit. In Table IV we have also given the values for these constants obtained previously ( 5 ) in 0.5M LiC104-DMS0 supporting electrolyte by potentiometric measurements with a thallium amalgam electrode. The agreement is good, considering the approximations inherent in our calculations from such sparse data and the different ionic media employed. Ion pairing of Li+ with C1- would be expected to be stronger than with T1Cl2-, and this difference would be reduced or effectively eliminated (Le., EtrN+ medium). Thus, Ks2is expected to be larger (log K,z more positive) in the Et4Nf medium than in the Lif medium, and this is observed. A similar argument implies that K,oshould be larger in the Lif medium, and this is observed, but the effect of higher ionic strength may offset the ion pairing effect somewhat. Ksl should be independent
of the ionic medium to a first approximation, but the discrepancies between the two results for this constant may be a reflection of experimental and theoretical errors. ACKNOWLEDGMENT
The authors thank Walter Zurosky for assisting with some of the experiments, P. A. Malachesky and M. Salomon for allowing us to quote their unpublished work, Susan Kirkland for distilling the propylene carbonate, and Raymond Jasinski and David Cogley for many helpful discussions. RECEIVED for review June 16, 1969. Accepted August 25, 1969. This work was sponsored by the Air Force Cambridge Research Laboratories, Office of Aerospace Research, under contract A F 19(628k6131, but does not necessarily constitute the opinion of that agency.
lodide-Selective Electrodes in Reacting and in Equilibrium Systems John H. Woodson' and Herman A. Liebhafsky2 Department of Chemistry, Texas A & M Unioersity, College Station, Texas 77843 WHENH 2 0 2is added to iodate in acid solution, iodine and reactive intermediate compounds-e.g., HIO, HI02, H2I203thereof are formed, a complex reaction system results, and hydrogen peroxide is catalytically decomposed ( I ) . Earlier investigations of this system (2) are being continued with major emphasis at present on measurements of the iodide activity, [I-], long recognized as a crucial variable and thought at times to be below 10-6M. The usual Ag/AgI electrode could not be trusted in this reactive system. Of much greater promise are the iodide-selectiveelectrodes (3) now available, in which crystalline silver iodide is interposed as a membrane between the unknown [I-] solution (here, the reaction mixture) and presumably an Ag/AgI electrode in contact with aqueous silver nitrate. When such an electrode is used in a reaction mixture, the membrane is a possible reactant. For example, if the mixture contains no silver ion, the reaction AgI (c, membrane)
=
Ag+ (as>
+ I- (as)
(1)
will proceed to establishment of the solubility equilibrium K?
AgI (c) e Ag+ (as)
+ I- ( a d
(2)
The iodide produced in Reaction 1 of course influences the electrode reading. When this iodide is at least comparable with the iodide to be measured, the electrode reading will be too large. It will approach %'%as a limit (the solubility limit of the electrode) as the initial [I-] approaches zero. Fluoride-selective electrodes in equlibrium systems (4, 5 ) 1
On leave from San Diego State College, San Diego, Calif. 92115.
* To whom correspondence should be directed.
(1) W. C . Bray and H. A. Liebhafsky, J. Amer. Chem. Sac., 53, 38 (1931). (2) H. A. Liebhafsky, unpublished work, University of California, Berkeley, 1928. (3) G. A. Rechnitz, Chem. Eng. News, 43, (25), 146 (1967). (4) M. S. Frant and J, W. Ross, Jr., Science, 154, 1553 (1966). (5) J. J. Lingane, ANAL.CHEhg., 40,935 (1968). 1894
and others (6) have been shown to behave in accord with the discussion just given. In reacting systems, fluoride-selective electrodes have been used successfully (7) to follow [F-] at activities (about l O - 4 M ) so far above the solubility limit that calibration presented no serious problems (8). In the present investigation, [I-] values near and below the solubility limit proved of concern. EXPERIMENTAL
Analytical Reagent or C.P. grade chemicals were used as purchased. The 30% H202 contained no preservative, An Orion iodide-selective electrode was used and its potential measured with an Orion 801 Digital pH/mV meter and recorded by a Hewlett-Packard 7100B recorder with a 17500A input module. Measured potentials were precise to 0.2 mV and recorded to 1 mV. The Ag,'AgI electrode was made by electrolyzing a silver billet electrode (L&N #117226) in 1M HI for several minutes at low (5-10 mA/cm2)current density and was aged several days before use. Its potential was measured directly with a 17500A input module on the same (two-channel) recorder. The reference electrode (Orion 90-02, double junction) contained a 1N NarS04 solution as salt bridge between the reacting system and the calomel reference electrode; this salt solution is compatible with the reacting system, which contained perchloric acid. In the longer experiments, some of which extended to several days, the sleeve-junction end of the reference electrode was covered with a collodion film (9) to retard flow of the bridging solution into the reaction mixture. The collodion had no significant disturbing effect. All [I-] measurements were performed in a borosilicate beaker at constant ( j ~ 0 . 0 2"C.) temperature in a water bath on solutions stirred magnetically with a Tefloncoated stirring bar. The rubber stopper closing the beaker and holding the electrodes was protected by a film of collodion to prevent serious attack by iodine vapor. The iodide(6) R. P. Buck, ANALCHEM., 40, 1432, 1439 (1968). (7) K. Srinivasian and G. A. Rechnitz, ibid., p 1955. (8) Zbid., p 509. (9) H. A. Fales and M. J. Stammelman, J. Amer. Chem. Soc., 95, 1272 (1923).
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13. NOVEMBER 1969
Table I. Thermodynamic Data for 298.15
AGO, (kcal/mole)
AgI(c) Ag+ (as) I- (as) A for equilibrium 2 a
-15.82 18.433 -12.23 21 ,923
OK5
c
AH"j
On
(kcal/mole) (cal/mole deg) -14.78 25.234 -13.19 26.82
13.58 5.2 -34.0 -42.38
See references I I and I2.
Table 11. Solubility Limits of Detection for Iodide-Selective and Ag/AgI Electrodes .i y
EMF
r74
.35
Figure 1. Nernstian response and solubility limits of an iodideselective electrode in various solutions of NaI at 50 "C. Iodide activity (PI = -log [I-] = -log y C,) DS. potential (EMF) 0.20MNaC104( p = 0.2, y = 0.72) W 0.067M &SO4 ( p 0.2, y = 0.72) A in water ( p = 0.00, y = 1.0) - Theoretical response in equilibriumsystems Extrapolation of Nernstian response (slope = 64.1 mV/pI unit) beyond the solubilitylimit - - - - Solubility limits (actual and theoretical) in NaCI04 solution
-
The EMF values for PI = m were arbitrarily plotted as shown to avoid crowding; the remaining points on each curve are plotted relative to this ordinate
selective electrode was mounted a t 30" from the vertical so that with rapid stirring no bubbles of oxygen were trapped on its recessed face. This measure eliminated wild voltage fluctuations caused by formation and detachment of gas bubbles on the sensing surface of a vertical electrode. When [I-] measurements were to be made on reacting systems, hydrogen peroxide was added to the rest of the reaction mixture. Oxygen evolution soon thereafter gave evidence that the peroxide was being catalytically decomposed by the 12-103- couple (10).
Equilibrium Systems. Measurements with the iodideselective electrode were made in known dilute solutions of NaI maintained at a constant ionic strength of 0.2 with either NaC104 or K2S04. The solid curve in Figure 1 shows the behavior to be expected on the basis of Reactions 1 and 2 when the iodideselective electrode described herein is used in a solution initially free of Ag+. Under these conditions, the Nernst equation reads as follows for the electromotive force produced by the difference in [I-] on the two sides of the membrane:
RT 5
+ y ) constant
1
Calculated (pK2/2) Manufacturer" Observed; in: NaCIOa(p = 0.2) KtSOa ( p = 0.2) (as) ( p = 0.0)
7.64 7.1
8.03 7.8
6.4 6.9 6.6
6.5
... ...
6.3 6.5 6.3
(3)
6.2 ... ...
a Instruction Manual Halide Ion Activity Electrodes, 2nd ed. (Form IM94-H3/767)Orion Research Inc., Cambridge, Mass. 1967
These conditions determine the solid curve in Figure 1. For x >> y , the usual Nernstian slope is to be expected; for y >> x, the solubility limit operates, and the electrode cannot be used to measure the initial [I-] in equilibrium systems to which no Ag+ is added. Equilibrium systems that act as iodide buffers a t low [I-], such as can be prepared from AgI(s), TII(s), and added Ag+, are outside the scope of this paper. For a general treatment that includes the case under consideration here see (6).
For K2 as a function of temperature over a limited range, chemical thermodynamics gives AGOTIT = - R In KZ = A H o I / T - ACop In T
RESULTS AND DISCUSSION
e = - In ( x
Limiting PI Iodide-selective Ag/AsI 50 "C 25 "C 50 "C 25 "C
+I
(6)
This equation has in effect been used to calculate K2 a t 50 "C from its value a t 25 "C; H o I and I are integration constants evaluated for the data in Table I. According to the foregoing discussion, [I-] values measured with the iodide-selective electrode should follow the slope predicted by Equation 3 when x >> y , and have the limiting value dk-when y >> x . The first requirement is met: slopes of 59.2 (25 "C) and 64.1 (50 "C) mV per unit PI were obtained. See Figure 1. The second requirement was met a t neither temperature although painstaking experiments (not all of which are reported) were done. See Table 11. Such behavior could have many and complex causes among which an ubiquitous iodide impurity and an increased solubility of the
where x is the initial [I-], which was to be measured, and y is the added [I-] ( = [Ag+]) contributed by Reaction 1. Evidently, ( x
+ y)y
and (for our conditions) y
5
=
KZ
fifor all x.
(10) J. H. Woodson and H. A. Liebhafsky, Nature, in press.
(4) (5)
(11) D. D. Wagmen, W. H. Evans, V. B. Parker, I. Halow, S . M. Bailey, and R. H. Schumm, National Bureau of Standards Technical Note 270-3, Table 12(1), U. S . Government Printing Office, Washington, D. C., Jan. 1968. (12) D. D. Wagmen, W. H. Evans, V. B. Parker, I. Halow, S. M. Bailey, and R. H. Schumm, National Bureau of Standards, Technical Note 270-4, Table 37(1), U. S . Government Printing Office, Washington, D. C., May 1969.
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
1895
-
I 1 '-
*L-.L 20
TIME
40
60
80
(MI%)
Figure 2. Smooth catalysis: Potential of an iodide-selective electrode us. time at 50 "C. Initial concentrations: 0.10M KIO3, 0.0765M HCIO4, 0.48M H20z. PI scale established by linear extrapolation of measured values for low PI. See Figure
. 1
membrane or of electrolytic silver iodide are intriguing possibilities. Because of discrepancies in Table 11, a second electrode was obtained and investigated with results identical to those reported in Figure 1 and Table 11. See Figure 4. It is certain, however, that a lower concentration limit for potentiometric determination of iodide ion exists in practice as well as theory for these electrodes in unbuffered iodide solutions. Reacting Systems. SMOOTH CATALYSIS. The PI values in Figure 2, obtained from recorded EMF values during the first 80 minutes after the addition of peroxide, are reasonable in terms of what is known for the experimental conditions about the reaction system from measurements of [Iz] and of oxygen evolution (2). Smooth catalytic decomposition of H20z into H 2 0 and (1/2) Oz begins after [I-] has reached a maximum near the close of the induction period. The proportionality of [I-] to [Hz02],which is indicated by the linear portion of the curve, parallels the previously observed proportionality between [I2]and [HzO2]. The important thing here is that the PI values in Figure 2 are wholly unaffected by the solubility limit of Figure 1. PERIODIC (PULSED)REACTION. The need to explain the periodic (pulse-like) characteristics of this reaction system is a major reason for resuming its investigation and for attempting to make [I-] measurements in it. That [I-] pulses occur synchronously with the pulses in oxygen-evolution rate was soon discovered and has been reported (11). Of present concern are three features shown in Figure 3: Striking pulses in [I-] are observed under the conditions given beneath the figure for both the iodide-selective electrode (upper half of figure) and the (conventional) Ag/AgI electrode; in the former case, the [I-] pulses oscillate undisturbed through the calculated solubility limit (Table 11) with remarkable regularity and lie entirely below the measured solubility limit (Figure 1); and the discrepancy between the pulses recorded by the two electrodes tends to disappear as the concentrations of reactive substances (H20z,Iz,and intermediate compounds) decreases. The iodide pulses show other features, notably the linear region when the minimum has been passed, that are undoubtedly of kinetic significance. Discussions of mechanism are, however, premature. Although no final judgment is yet possible about the kinetic significanceof the [I-] measurements, it is difficult to avoid the conclusion that they must at least be relatively significant. If so, then it is reasonable to explain as follows why the sohbility limit is meaningless in this reacting system. Some of the participating reactions remove iodide so rapidly-including 1896
.*53,
1
9C
I
Av
90
I 28s
1
I
25-
a
-
TlKE r * n i
Figure 3. Pulsed catalysis : Potential of an iodide-selective electrode (upper trace) and potential of an Ag/AgI electrode (lower trace) us. time at 50 "C. Initial concentrations: 0.10M K103, 0.056M HC104, 0.48M HsOe. PI scale established by linear extrapolation of measured values at low PI. See Figure 1
iodide formed in Reaction I-that the [I-] in the system is governed primarily by kinetic considerations: Reaction 1 may thus provide the [Agf] required to establish Equilibrium 2 without raising the [I-] above its kinetically controlled value. It was noticed that after lengthy exposure to the reacting system a powdery yellow film was deposited on the sensing end of the electrode, both on the membrane and on the adjacent plastic body of the electrode. If the film is allowed to accumulate, the potential readings eventually become unsteady and show random fluctuations as large as 15 mV. Gentle polishing of the membrane with a soft paper tissue removed this film and restored the performance of the electrode. The iodide-selectiveelectrode was compared to the conven-
0
e 0
.*a
EMF
0 0
D 3
n
r
*49
+4c:
J
Figure 4. Nernstian response and solubility limits of several electrodes in various solutions of NaI and 0.20M NaC104 at 50 "C. Iodide activity (PI = -logy CI; p = 0.20, y = 0.72) us. potential (EMF) 0
v
Iodide-selective electrode #l(same as Figure 1) Iodide-selectiveelectrode #2 -. - electrode
- Theoretical Nernsiian response, slope
The EMF values for avoid crowding; the relative to this ordinate
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
=
= 64.1 mV/pI unit. were arbitrarily plotted as to points on each are plotted
tional Ag/AgI electrode in equilibrium systems also (see Figure 4). The following observations are in order: The solubility limit is somewhat higher for the Ag/AgI electrode; the response of this electrode to changes in [I-] was more sluggish, perhaps because equilibrium established itself more slowly on the finely divided AgI(s); and in an acidified iodide solution through which oxygen is bubbled, the potential of the Ag/AgI electrode is shifted to more positive values-i.e., in the direction that would increase the difference between the potentials of the two electrodes in Figure 3. There seems
little doubt that the presence of oxidizing agents produces a mixed potential at a n Ag/AgI electrode. ACKNOWLEDGMENT.
The many helpful suggestions of Karl H. Pearson are gratefully acknowledged. RECEIVED for review June 30, 1969. Accepted August 27, 1969. This work was supported by the Robert A. Welch Foundation under Grant A-254.
Sensitized Cation Selective Electrode J. Montalvo, Jr., and G . G . Guilbault Department of Chemistry, Louisiana State University in New Orleans, New Orleans, La. 70122
IN A RECENT STUDY (I), the results of a thorough investigation of the response, selectivity and use of the cation electrode for determination of NH4+was reported. The application of this electrode to a study of the kinetics of deaminase enzyme systems (urease, asparaginase, glutaminase, amino acid oxidase, and amine oxidases) was likewise reported. In a more recent study (2), urea was determined by immobilizing the enzyme urease in a layer of acrylamide polymer on the glass surface of a cation electrode sensitive to ammonium ion. The substrate urea diffuses to the enzyme electrode and reacts with the immobilized enzyme to produce ammonium ion at the surface of the electrode. In this paper, the use of a film of immobilized urease on a glass electrode is reported
for the determination of cations. EXPERIMENTAL
Preparation of the Coated Electrode. The gel solution was prepared by dissolving 3.0 grams of acrylamide monomer and 0.58 gram of N,N’-methylene-bisacrylamide (Eastman) in 25 ml of 0.1M tris buffer, pH 7.0. To catalyze the photopolymerization, 2.7 mg of riboflavin and 2.7 mg of potassium persulfate were added. The solution was stored in the dark at room temperature until used and remade every two days. One milliliter of the polymer solution was pipetted into a small centrifuge tube containing a weighed amount of enzyme, 175 mg urease/cc of gel solution. The enzyme suspension was stirred for two minutes and was allowed to dissolve for 20 min at room temperature. The mixture was chilled in the refrigerator at 2 “ C for 10 min and centrifuged for 20 min. The clear supernatent was carefully transferred to a test tube with a 21-mrn i.d. A monovalent cation electrode (Beckman Instrument Co., Catalog Number 39137) was washed well with distilled water, wiped dry with tissue paper, and mounted upside down. A strip of nylon net, 350 p thick, (nylon stocking, J. C. Penney and Co.), 2 ” X 2”, was placed over the glass bulb of the electrode. A rubber “0” ring was used to anchor the netting over the rigid glass bulb. This “0’ring was placed about inch below the sensing part of the glass bulb and held the netting in place while a second “0” ring was (1) G. G. Guilbault, R. Smith, and J. Montalvo, Jr., ANAL.CHEM., 41, 600 (1969). (2) G. G. Guilbault and J. Montalvo, Jr., J . Amer. Chem. SOC., 91, 2164 (1969).
placed just below the edge of the sensing glass bulb. The first “0” ring was cut away and the netting over the glass bulb carefully inspected for any folds which can be removed by pulling on the netting below the remaining “0” ring. The netting was then cut flush with the “0” ring. The resulting electrode was dipped into the enzyme gel solution making sure all of the pores of the netting are filled with solution. The electrode was then removed from the solution and the glass wall above the “0” ring wiped free of solution. Excess liquid does not cling to the netting because of the surface active enzyme preparation. The netting gives mechanical rigidity to the enzyme gel layer and also is used to control its thickness. The electrode was placed in a water-jacketed tube. Oxygen inhibits the polymerization and was removed by bubbling with NP through the tube for 15 min. The polymerization reaction which requires light, was irradiated with a G. E. BBA photoflood lamp equipped with a reflector. The controlled temperature in the photopolymerization was 28 “C (measured with a mercury bulb thermometer in place of the electrode). After one hour of polymerization, the electrode was equilibrated in tris buffer for one day before use. Apparatus. A Beckman Zeromatic Model I1 pH meter and a standard fiberjunction, saturated calomel reference electrode (SCE) were used. Millivolt measurements were made by operating the pH meter in the h 7 0 0 mV range and were recorded on an Electroscan 30 (chart speed of the recorder varied from 20 to 100 seciinch). In those studies involving Ag+ ion a salt bridge filled with tris-HNOs buffer was used to prevent contamination of the SCE. All measurements were carried out in a thermostated cell at 25 f 0.01 OC. Chemicals. Tris(hydroxymethy1amino-methane)buffer, was used to maintain a constant pH 7.0, and a constant ionic strength, 0.1M. The enzyme urease was obtained from the California, Corp. for Biochemical Research, Grade B (activity 375 Sumner units per gram of enzyme). Procedure for Determination of Cations. Potentiometric measiirements of the steady state response for the construction of calibration curves and the study of the effect of various parameters (film thickness, gel concentration, etc.) on the steady state response for the coated electrode were carried out in the conventional manner. All solutions were magnetically stirred with a bar made of Teflon (Du Pont). An aliquot of the cation to be determined (NH4+, Na+, etc.) is pipetted into a 100-ml beaker containing 50-ml of buffer. The steady state potential is read, and the ion concentration of cation is determined from a calibration plot of potential us. log of cation concentration. After determination of an ion, the film of immobilized enzyme was washed free of the ion by placing both the coated electrode and the reference
ANALYTlCAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
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