+
Table 11. Comparison of Concentration Ranges in the SDS-CPC Titration Used by Different Workers
Original concn of SDS solution, mM/ml
Authors Epton ( I ) 3.00-8.00 Weatherburn (2) 1 . 7 4 Present work 3.47-13.88
Final volume of aqueous phase after Volume of SDS addition solution of Me B solution, used, ml ml 10 1C-40 60
35 75 85
Final concn of SDS in aqueous phase, mM/ml 0.86-2.29 0.23-0.93 2.45-9.80
Schick and Fowkes (9) found that the cmc of SDS was reduced to one third of its initial value in the presence of 22.2 lauryl alcohol, the present solutions containing initially 1 mg of SDS/ml (3.47mM/ml) would still be expected to be below the cmc (0.238 wt. (8.lmM) for pure aqueous SDS) based on the volume of the aqueous phase after addition of CPC solution. The presence of micelles in the solution apparently does not interfere with the exchange reaction, (9) M. J. Schick and F. M. Fowkes, J . Phys. Chem., 61,1062 (1957).
+
SDS.methylene blue CPC + SDS-CPC methylene blue, which provides the end point of the titration by developing a blue color in the aqueous phase. Two additional conclusions can be drawn from the data in Table I. Although variable times of standing between 10 and 30 minutes were allowed after addition of lauryl alcohol before titration with CPC, good precision among replicate runs was obtained in all cases. This suggests that the solubilization of lauryl alcohol by the SDS solution was complete in 10 minutes, although there has been no study of the kinetics of this process. The absence of a dilution effect, as reported by Epton ( I ) , in the present results may be due to use of a higher concentration range of SDS (cf. Table 11). This would be in accord with Epton’s observation that the dilution effect became smaller at higher concentrations. A comparison of the concentration ranges of SDS studied by different workers is shown in Table 11. It is evident from these results that the method is satisfactorily applicable over a wide range of concentration of SDS from 0.86 to 9.80mM, (based on the concentration in the aqueous phase after addition of methylene blue solution but before titration with CPC solution), and need not be restricted to the concentration range between 0.003 and 0.008M used by Epton (1). RECEIVED for review July 22, 1971. Accepted November 4, 1971. Support of a fellowship which made possible this work is gratefully acknowledged to the Foods Division of the Anderson Clayton Company.
Use of Metal Tungsten Bronze Electrodes in Chemical Analysis M. A. Wechter,’ H. R. Shanks, G. Carter, G. M. Ebert, R. Guglielmino, and A. F. Voigt Institute for Atomic Research and Departments of Chemistry and Physics, Iowa State University, Ames, Iowa
THEMETAL TUNGSTEN BRONZES, nonstoichiometric compounds of formula M,W03 function as indicating electrodes in several widely differing electrochemical systems. Current investigations indicate that electrodes made of the bronzes may be used to determine pH, pMetal for certain reducible species such as Ag(1) and Hg(II), and to follow potential changes involved in some oxidation-reduction systems. Even more significant perhaps, is that they appear to function as oxygen sensitive electrodes. The use of tungsten bronzes as “catalytic” electrodes in fuel cells has been reported ( I - j ) , but little has been published on applications as measuring electrodes. Recently a Russian group has published several papers (4-6) on uses of bronzes 1 Present address, Chemistry Department, Purdue UniversityCalumet, Hammond, Ind.
(1) A. Damjanovic, D. Sepa, and J. O’M. Bockris, J. Res. Inst. Catal. Hokkaido Univ., 16, 1 (1968). f2) D. Seva. A. Damianovic, and J. O’M. Bockris, Electrochim. Acta, 12,746 (1967): 13) L. W. Niedrach and H. I. Zeliger, - J. Electrochem. SOC.,116, 152 (1969). (4) A. G . Kohsharov and V. F. Ust-Kachkintsev, Uch. Zap. Perm. Gos. Univ., 111,63 (1964); Chem. Abstr., 64,16178(1966). (5) Ibid, 178, 117 (1968); Chem. Abstr., 73,6197e (1970).
.,
850
ANALYTICAL CHEMISTRY, VOL. 44, NO. 4, APRIL 1972
in pH and other measurements. At this time details of these experiments are not available to the authors. The metal tungsten bronze crystals used for this work were prepared by the fused salt electrolysis of the appropriate metal tungstate and WOs. The details of the single crystal growth are available elsewhere (7). The compositions of the bronze crystals were determined either directly by neutron activation analysis or from the measurements of lattice constants. The relationship between lattice constants and composition has been determined by neutron activation analysis (8). Individual pieces, single crystals, of bronze to be used for electrodes were cut from larger crystals with a diamond saw. The electrodes were prepared by attaching the crystal to one end of a glass tube with epoxy cement. The tube was partly filled with mercury which made electrical contact to the crystal. A copper wire could then be inserted in the mercury and used as an electrical lead to the electrometer used for potential measurements. Details of the electrode are shown in Figure 1. (6) A. G. Koksharov and V. F. Ust-Kachkintsev, lzv. Vyssh. Ucheb. Zaved. Khim. Khim. Tekhnol., 10, 243 (1967); Chem. Abstr., 67, 39.5126 (1967). (7) H. R. Shanks, J. Crystal Growth, in press. (8) M. A Wechter, H. R. Shanks, and A. F. Voigt, Inorg Chem., 7, 845 (1968).
Ill
Figure 1.
Ill
The metal bronze electrode
Dimensions: tubing 12 cm X 0.6 cm 0.d. crystal 0.7 X 0.08 X 0.15 cm
0
I
I
I
I
I
I
I
I
5
I 7
6
I
8
I 9
I IO
I
I
I
I
I1
12
I3
14
15
ml HCIO4 -100
Figure 3. Titration of 1.5 X 10-2MC6H6NH1with 0.1M NaOH in glacial acetic acid
-200
450
I
I
I
2.0
3.0
4.0
I
I
5.0
6.0
400
-2 -
-300
J
2zw
i:I
-400
5
-500
IO0
50
-600
0
2
4
6
10
0
I2
14
16
18
ml NoOH
Figure 2. Titration of 3 NaOH
x
10-8M HC104 with 0.1M
EXPERIMENTAL AND RESULTS
Preliminary work has yielded the following results. pH Measurements. A number of titrations have been performed both in aqueous and in non-aqueous media. The electrode pair consisted of a sodium tungsten bronze with x = 0.7122 which functioned as the indicating electrode, and a saturated calomel electrode which served as a reference. Data for some typical titrations are plotted in Figures 2 and 3. Figure 2 represents the titration of -400 ml of 3 X 10-3M HC104 with 0.1M NaOH. The curve obtained is a normal titration curve, exhibiting the large break at the end point which is expected in strong acid US. strong base systems. Figure 3 represents the non-aqueous titration of -100 ml of 1.5 X 10-2M aniline with 0.1M HClOd in glacial acetic acid. The plot of potential us. milliliters of titrant is typical for the system. From these and other data, it first appeared that the tungsten bronze electrode functioned remarkably like the glass electrode for pH measurements. However, the bronze and glass electrodes were compared by immersing each in turn in a series of standard buffer solutions, and obtaining for each system a plot of potential us. pH. The saturated calomel electrode was, in each case, used as the reference electrode. The glass electrode system yielded the expected straight line graph, with a slope of approximately
t
1.00
7.0
PA9
Figure 4. Plot of potential (mV) us. pAg
59 mV/pH unit. On the other hand, the graph of the bronze system was curved, indicating that, under present conditions at least, the pH response of the bronze electrodes deviates from linearity. Several titrations were performed using both bronze and glass systems in solution. In each case HC104 was titrated with NaOH. The titration curves for the bronze-calomel system closely resembled those for the glass-calomel pair, except that the break at the end point for the former pair appeared to be 10-50z larger than that for the latter. It is apparent that, while the tungsten bronze electrodes may not be presently useful in the direct determination of pH, they can be used to follow pH titrations. pMetal Measurements. Certain metal ions in solution will elicit responses from the bronze electrodes. From preliminary studies, it appears that only those metal ions whose reduction potentials are positive with respect to the H+, 1/2H2couple will cause the potential of these electrodes to vary. No response was observed on the addition of metal ions of Groups IA, IIA, IIIA, and others whose reduction potentials are negative. However, for ions such as Ag(I), Hg(II), Fe(III), and Cu(II), a change in concentration results in a change in electrode potential. Figures 4 and 5 represent data taken for Ag(1) and Hg(I1). For Ag(1) (Figure 4), there is a nearly linear relationship between pAg and potential response. This experiment was performed in 0.01N HzS04, with saturated calomel as the reference electrode. ConANALYTICAL CHEMISTRY, VOL. 44, NO. 4, APRIL 1972
*
851
tioo
The low solubility of HgS04 eliminated the possible use of a i sulfate medium to avoid this effect. From experimental results obtained thus far, in particular
400
->
300
+
I
-I
?E
200
?i
loot I
I
I
I
I
I
1.0
2.0
3.0
4.0
5.0
6.0
\ I
7.0
8.0
P Ha
Figure 5. Plot of potential (mV) us. pHg
0
BRONZE VS. S.C.E.
X
PI VS S.C.E.
A BRONZE M Pt
I
I900 - f
0
2
4
6
8
10
12 w ml Ce tpI)
i
i
I1
I 13
l I5
l I7
I9
Figure 6 . Redox titration curves centration of the silver ion was increased regularly, and potential measurements were taken l minute following each addition of silver. In order to avoid precipitation of AgCl which might result because of the contribution of chloride ion from the calomel electrode, contact was made between the reference electrode and the solution by means of a salt bridge. In Figure 4 the slope of the linear portion of the potential us. pAg graph approximates 59 mV/pAg unit, the value expected from the Nernst equation for a one electron change. Other curves for the Ag(1) system were obtained using 0.01M NaNO, as electrolyte. These curves showed a substantial break or plateau at pAg 3-4, which corresponds to the precipitation of chloride ion present as an impurity in the NaNO,. Obviously, then, this method could also be applied to the determination of C1- and perhaps other anions as well. Figure 5 is a plot of potential us. Hg(I1) concentration with data obtained as in the case of Ag(1) except that the electroThe slope of the curve at pHg lyte was 0.01M "0,. below 4.3 and above 5.7 was approximately 30 mV/pHg unit indicating a two-electron change. As in the Ag experiments in nitrate solution, a break in the curve was observed, in this case a considerable change in slope between pHg 4.3 and 5.7. It appears likely that this break has an explanation similar to that used for the Ag-NaNO, systemnamely, the formation of a precipitate or complex with the traces of chloride or other impurity in the nitrate medium. 852
visual observations of the electrodes after use, it is proposed that the response mechanism is the reduction of the metal ion at the surface of the electrode. From a consideration of the high conductivity, 3.5 X l o 4 ohm-' cm-1, of the tungsten bronzes in this composition range, this seems to be a reasonable mechanism. Redox Titrations. Preliminary work based on the familiar oxidation-reduction reaction Fe(I1) Ce(1V) = Fe(II1) Ce(II1) shows the utility of the tungsten bronze electrodes for redox titrations. All solutions were prepared in 0.5M H2S04 to ensure constant pH during titration. Three separate titration curves were obtained for each determination, and a representative series appears in Figure 6. The electrode systems used were 1) bronze us. calomel, 2) Pt us. calomel, and 3) bronze us. Pt. Electrode systems 1 and 2 yielded normal titration curves, with system 1 lagging system 2 by an increment which was maximized in the vicinity of the end point. System 3 represents the difference between the first 2 systems and takes the shape of a first derivative curve. It appears likely that the bronze us. Pt electrode system will also function in other redox systems as a bimetallic electrode. The advantages of obtaining first derivative curves over normal curves are obvious, particularly in application to automated titration devices. Oxygen Measurements. Results obtained from pH measurements led to the idea that oxygen may have a measurable effect on the potential. Investigations to date show that the bronze electrodes are, in fact quite sensitive to dissolved oxygen, particularly at high pH. Accurate independent measurements of pOz at high values are difficult, but comparison with polarographic determination of pOa indicates that at pH 12, AE/ApO?is of the order of 100 to 120 mV/pOz unit (9).
ANALYTICAL CHEMISTRY, VOL. 44, NO. 4, A P R I L 1972
+
CONCLUSIONS
The versatility of the metal tungsten bronzes for use as electrodes has been demonstrated. Sodium tungsten bronzes, Na,W03, were used in this work because they were readily available as large single crystals. However, measurements made with other tungsten bronzes, specifically Li,W03 and Rb,W03, have yielded similar results. The bronze electrodes can be used to great advantage in measuring pH, pMetal, and redox responses, and an outstanding application appears to be in the measurement of ~ O Z . A clean reproducible surface must be maintained, as would be expected if the postulated mechanism, reduction of the species at the electrode surface, is correct. The clean surface can be obtained with a base, such as NH3 or NaOH, a treatment which was used quite effectively in work with the silver system. The surface can also be cleaned and regenerated with an air-sand abrasive. On the basis of the proposed mechanism, other applications suggest themselves, such as the use of powdered bronzes as reductors for reducible organic as well as inorganic species. Experiments in progress are based primarily on pH, pM, pOz, and redox responses. Future work will include: function of powdered bronzes as reductors; study of the electrode response in complexometric titrations ; bronze electrode, radio-tracer studies on interferences in pMetal determinations, and for more clearly defining the reaction mech(9) M. A. Wechter, P. B. Hahn, D. C . Johnson, and A. F. Voigt, Iowa State University, Ames, Iowa, unpublished work, 1971.
anism; variations in electrode response with variations in bronze composition; and kinetics involved in redox systems, using the bronze electrodes and radio-tracer techniques.
are also grateful for discussions with Paul Hahn, Ames Laboratory, and Harvey Diehl and Dennis Johnson of the Chemistry Department, Iowa State University.
ACKNOWLEDGMENT
The authors acknowledge Gordon Danielson, Ames Laboratory, for providing the tungsten bronzes, and Paul Millis, Ames Laboratory, for his technical assistance. We
RECEIVED For review August 20, 1971. Accepted October 26, 1971. Work performed in the Ames Laboratory of the U.S. Atomic Energy Commission.
Role of Solvent Extraction Parameters in Liquid Membrane Ion Sektive Electrodes Helen J. James, Gary P. Carmack, and Henry Freiser Department of Chemistry, University of Arizona, Tucson, Ariz. 85721
THERELATIVELY RECENT development of a new class of electrochemical sensors called ion selective electrodes has quite properly been the subject of widespread attention and study ( I ) . Of these, the type referred to as liquid membrane electrodes offer the exciting prospect that solvent extraction principles provide a guide to their design and operation. Thus, most of the liquid membrane electrodes introduced so far incorporate solvent extraction systems. Metal ion electrodes have made use of their extractable complexes with acid phosphate esters and thiocarboxylic acids. Anion responsive electrodes have employed extractable ion association complexes using ferroin and analogous large cations (2). Recently, we have described a series of anion responsive electrodes employing the extractable complexes formed by the tricaprylmethylammonium cation (Aliquat 3363) and a number of inorganic and organic cations (3-5), in which the qualitative observation of the relation of electrode selectivity to ion pair extractability was made. In this communication we wish to report the results of a quantitative study of the extraction equilibria of a representative series of ion pair complexes undertaken to examine the postulated relationship more closely.
culture tubes (15 X 125 mm), covered with Parafilm, were used as counting containers. Extraction Equilibrium Experiments. DISTRIBUTION OF THAI. Equal volumes of a 3.30 X 10-3M THAI solution in an organic solvent and an aqueous phase containing a sufficient amount of 13II to give good counting statistics were shaken together for 30 minutes, a time adequate for equilibration. The phases were allowed to separate and 5 ml aliquots of each phase were pipetted into separate counting tubes and counted. The ratio of activities observed was taken as a measure of D, the distribution ratio. In order to avoid possible interferences from trace impurities, the organic phase was separated, equilibrated again with distilled water, and the value of D redetermined radiochemically. Agreement was obtained to within lO-15z (0.05-0.07 log unit). COMPETITIVE EXTRACTION EQUILIBRIA.Equal volumes of an aqueous solution containing 3.30 X 10-2M organic anion and a sufficient amount of 1311to give good counting statistics and an organic solution of 3.30 X 10-3M THAI were equilibrated and aliquots of the separated phase counted as previously described. With the amino acid distribution series, the aqueous phases were adjusted to pH 10.5 with NaOH prior to equilibration. The equilibrium pH values of the amino acid solutions were also determined. Agreement in these experiments was 5 or better.
EXPERIMENTAL
Reagents. Tetra-n-hexylammonium iodide (THAI) was obtained from Eastman Organic and used as received. All other reagents used in this study were of analytical reagent grade. Distribution studies made use of carrier-free lalI (New England Nuclear Corp.) in the form of NaI. Apparatus. Extractions were performed using 45-ml cylindrical glass vials fitted with polyethylene thimble stoppers and plastic screw caps. Samples were shaken in an Eberbach reciprocating shaker at the high speed setting with temperature control being maintained at 25 f 0.2 “C by circulating water from a Wilkens-Anderson Co. Lo-Temp bath through the jacketed shaker tray. A Nuclear-Chicago Model 186 scaler in conjunction with a Nuclear-Chicago Model DS-55 well-type scintillation detector was employed for radioisotope counting. Kimax lipless (1) “Ion-Selective Electrodes,” R. A. Durst, Ed., Nut. Bur. Stand. (US.)Spec. Publ. 314, Washington, D.C., Nov. 1969. (2) J. W. Ross, ibid.. Chapter 2. (3) C. J. Coetzee and H. Freiser, ANAL.CHEM., 40, 2071 (1968). (4) Zbid.,41, 1128 (1969). ( 5 ) M. Matsui and H. Freiser, Anal. Lett., 3, 161 (1970).
RESULTS AND DISCUSSION
It is relatively simple and convenient to quantitatively determine the participating equilibria of an ion pair complex extraction system containing an easily measured ion such as I3lI. Furthermore, by means of competitive extraction (i.e., a reaction system containing both 1311and another anion capable of forming an extractable ion pair complex), the convenience of radiochemical analysis serves for the determination of the equilibria involved in a whole series of extractable ion pair complexes. The same competitive reaction principle undoubtedly would apply to a series of complexes in which one of the ions has a readily measured characteristic ( e . g . ,high molar absorptivity, sensitive atomic absorption line, etc.). The distribution of THAI between an organic solvent and water involve two equilibria: ion pair formation in the aqueous phase Q+
+ I- =KIP +=
Q+, I--
ANALYTICAL CHEMISTRY, VOL. 44, NO. 4, APRIL 1972
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