would be an advantage in trace analysis where supporting electrolyte purity is often a problem. Since iR drop is proportional to the product of solution resistivity and drop radius (Equation 3), the iR drop could be maintained a t an acceptably low level for solutions of much higher resistivity by using drops of smaller radius than those used in the present work. The advantage of being able to record ac polarograms
without the need for positive feedback iR compensating circuitry, coupled with the fast scan rates of potential used (8), makes the rapid phase-sensitive technique attractive for routine analysis in nonaqueous or other high resistance solvents. Received for review June 29, 1973. Accepted December 10, 1973.
Potentiometric Microtitration of Sulfate Ion Using a Sodium-Selective Glass Electrode in a Nonaqueous Medium Naoshige Akirnoto Faculty of Pharmaceutical Sciences, Kyoto University, Kyoto 606, Japan
Keiichiro Hozurni Kyoto College of Pharmacy, Kyoto 607, Japan
The potentiometric titration of sulfate with an ion-selective electrode has been carried out using a lead-selective electrode to determine the end point ( I , 2 ) . However, the method is not acceptable for a highly accurate titration of sulfate in any low concentration, because the electrode does not give a curve with a well defined point of inflection that can be related to the equivalence point. On the other hand, the NAS 11-18 and NAS 27-4 (Corning Glass Works, Medfield, Mass.) have been used as representative glass electrodes which respond with high selectivity to monovalent cations. The former is known to have a selectivity order of hydrogen, sodium, and potassium ion; the latter, a reverse order of potassium, sodium, and hydrogen ion (3, 4 ) . They are commercially available under the respective specifications "Sodium-Selective Glass Electrode" and "Monovalent Cation-Selective Glass Electrode," in keeping with their most common application. Their response characteristics for monovalent cations have been examined in detail, but to date very little information has been reported concerning their response to divalent cations. In a study of this latter property, the present authors have discovered a potential break for barium ion when the electrodes were used in sulfate titration. The NAS 11-18 electrode in particular exhibits a sharp potential break. Subsequent investigation with regard to its application to the microdetermination of sulfate has disclosed that it will register a sudden increase in electrode potential in the vicinity of the equivalence point ( 5 ) under the condition of about a 70 vol 9i concentration of some organic solvent a t pH 5-6. However, the measurement of the Nernstian slope by plotting steady-state potentials in nonaqueous solutions of various concentrations of barium ion indicates no electrode response to barium ion, whereas the electrode, when exposed to a sudden change of barium ion concentration, developed a potential instantly which then gradually (1) J. W. Ross, J r . , and M . S. Frant, Anal. Chem.. 41, 967 (1969). ( 2 ) W . Selig, Mikrochim. Acta. 1970, 168. (3) G . Eisenrnan, "Electrochemistry of Cation-Sensitive Glass Electrodes" in "Advances in Analytical Chemistry and Instromentation," C. N . Reilley, Ed., Vol. 4 , Wiley (Interscience), New Y o r k , N . Y . . 1965, pp 213-369. (4) G . Eisenrnan, "Glass Electrodes for Hydrogen and Other Cations," Marcel Dekker. New York. N . Y . , 1967. (5) K . Hozumi and N . Akimoto, Anal. Chem.. 42, 1312 (1970).
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dropped off to the initial potential. This phenomenon is similar to the transient response to calcium and strontium ions reported by Rechnitz et al. ( 4 , 6, 7). One must therefore conclude that the potential break in the vicinity of the equivalence point during the titration of sulfate is due to a transient response of the electrode to barium ion. EXPERIMENTAL Apparatus. Automatic titrations were carried out with a Metrohm Potentiograph E 436 (The Metrohm Co., Herisau, Switzerland) set to the 500-mV full-scale range. The glass electrodes used were Corning NAS 11-18 (Cat. No. 476210) and Corning NAS 27-4 (Cat. No. 476220). The double-junction reference electrode was the Orion Model 90-02 (Orion Research Inc., Cambridge, Mass.) with a bridge solution of 0.2M hexamethylenetetramine ("Hexamine") to prevent any diffusion of potassium ion from the electrode to the sample solution. Reagents. A standard solution of 0.005M barium chloride was prepared by dissolving 122 mg of barium chloride dihydrate in 300 ml of water and diluting it t o a liter with isopropanol. A 0.2M Hexamine solution was prepared by dissolving 28.08 grams of Hexamine in water and diluting up to a liter. Procedure. An approximate volume of sample solution equivalent to 0.5 to 5.0 mg of sulfate radical (about 5 ml) is accurately pipetted into a 100-ml beaker and the solution is adjusted to pH 5-6 by adding 1 ml of the 0.2M Hexamine. Acetone and distilled water are added to bring the volume up to 50 ml and a t the same time to make it approximately 70 vol 9'0 in acetone. Both the sensing electrode and the reference electrode are rinsed well with 0.01M hydrochloric acid before immersion in the sample solution. Automatic titration with standard 0.005M barium chloride is then started at a delivery speed of 0.5 ml per min. The equivalence point in the titration is indicated by a sharp point of inflection in the titration curve on the recorder chart. In the event that the sample solution should contain other metal cations besides barium to which the electrode will respond, it is advisable that they be removed prior to the titration by passage of the solution through an ion exchange column. Otherwise, low analytical figures will be obtained.
RESULTS AND DISCUSSION Titration Curves in Highly Nonaqueous Medium. Titration curves of 2.500 ml of 0.005M sulfate obtained with the NAS 11-18 and NAS 27-4 electrodes in 90 vol 70 ace(6) G . A. Rechnitz and G . C. Kugler. Anal. Chem.. 39, 1682 (1967). (7) "Ion-Selective Electrodes," R . A . Durst. E d . . Nat. Bur. Stand. (U.S.).Spec.Pub/., 314, 313 (1969).
-150
I
2 -170 A
9 -125
? 2 A
II
-1501 -170
~
-150 -170
1I
-175
I 1
3
25
2
15
150-170
1
1
15
2
25
V
I
15
2
25
3
35
0.005M B a C I 2 , rnl
V
05
IV
-170
35
-200
I
r
-150
0 0 0 5 M BaC12, ml
0
111
3
35
0.005 M E a C 1 2 , ml
Figure 2. Effect of p H on titration of 2.500 ml of 0.005M sulfate in 90 vol % acetone I: pH 2; II: pH 3: I l l : pH 4: I V : pH 5: V: pH 6
tone a t p H 5.5 are shown in Figure 1. The potential jump produces the asymmetrical S-shaped curve which is characteristic of potentiometric titrations with ion-selective electrodes. In each case, however, the potential jump was not especially high, although a somewhat sharper break was observed with the NAS 11-18 electrode. Since this phenomenon immediately suggested the possibility of using the electrode for the microdetermination of sulfate, an investigation was undertaken to improve the titration as much as possible in the application of the NAS 11-18 electrode. Effect of pH. Titration data in 90 vol % acetone a t various pH levels obtained by adjustment with hydrochloric acid or with Hexamine (8) are shown in Figure 2. Although the potential break a t the equivalence point was not evident a t p H 2, it became progressively more apparent with increasing pH. Above p H 5, however, there was no further improvement. But if the titration was carried out a t p H 7 or above with the aid of a tris-buffer instead of Hexamine, the point of inflection of the titration curve shifted slightly in the direction of a slightly larger volume of titrant. Consequently, it is suggested that the p H of the sample solution be adjusted to p H 5-6 with a 1-ml volume of 0.2M Hexamine. Organic Solvent System. Titrations a t pH 5.5 in the presence of various organic solvents-ie., methanol, ethanol, isopropanol, and acetone-in 90 vol % concentrations were carried out in order to study the effect of each on the titration curve. Among these, methanol yielded a smaller potential jump than any of the others tried and the break came a little too soon, while the other three solvents yielded about the same titration curves. Nevertheless, ac( 8 ) H Shimada K Ono, and K Kawada, Ann Sankyo Res Lab (1969)
21, 40
etone was still recommended as the best solvent to use because it gave a slightly sharper and greater potential break than either ethanol or isopropanol. Figure 3 shows a series of titrations a t different acetone concentrations (pH 5.5). An asymmetrical S-shaped curve appears for a 90 vol YO concentration. With decreasing acetone concentration, the potential jump becomes less evident, but the electrode potential always remains steady u p to the equivalence point. Here, there is always the same remarkably sharp bend on the recorder chart. This is quite obvious in 70 vol % acetone, but not so striking a t 50 vol % where the potential jump is low. Therefore, an approximate 70 vol YO is recommended as a standard titration condition for the moment. To this end, the 0.005M barium chloride titrant should be made up in 70 vol 7'0 acetone in order to keep the titration solution at the same level in acetone throughout the whole titration. Delivery Speed of Titration. Repeated titrations of 15 samples of 2.500-ml standard sulfuric acid with 0.005M barium chloride were made at delivery speeds of 0.25, 0.5, and 1 ml/min, and standard deviations of 4, 4, and 7 (pl) as the measure of the reproducibility were obtained, respectively. Delivery speed of 0.5 ml/min is therefore advised for the titration, although 1 ml/min is still acceptable for ordinary microanalytical work. When successive titrations were carried out with water as the wash liquid for the electrode before each new run, some hysteresis was evident in connection with the electrode potential-ie., the potential began to change just short of the end point, thus giving rise to a titration curve with a dull point of inflection. This malfunction can be eliminated, however, by rinsing off the electrode with 0.01M hydrochloric acid each time before titrating. A hypothetical explanation for the difficulty experienced when washing with water may lie in the possibility of a certain amount of ion exchange between barium and hydrogen ions at the membrane surface of the sensing electrode after repeated use in an almost neutral medium. Washing with hydrochloric acid, on the other hand. may take off the barium and replenish the surface with hydrogen ion. Interpretation of Titration Curve. Ordinary glass electrodes for pH work show no potential break at the equivalence point in the titration of sulfate with barium ion in contrast to glass electrodes which are highly responsive to monovalent cations. In addition, the titrations described were carried out in a buffered medium through the use of Hexamine. These facts preclude change of hydrogen ion concentration during the titration. On the other hand, it is also recognized that the NAS 11-18 electrode is almost nonresponsive to divalent cations on account of its very low selectivity. But reports on its behavior in nonaqueous media have not been published to A N A L Y T I C A L C H E M I S T R Y , VOL. 46, NO. 6, M A Y 1974
767
‘A
2 I
-100O :
-200
1 ’ 1
7
-1 0
I
I C
6
5
4
3
2
0
1
10
20
Figure 4. Electrode responses of NAS 11-18 glass electrode to Na+, H + , K + , and Ba2+ concentrations in 90 vol % acetone
medium
30
TIME, min
-LOG CCATIONI
Transient response of N A S 11-18 glass electrode to a sudden change of barium ion concentration in 90 vol ‘YOacetone medium at pH 6.5
Figure 5.
Solutions containing metal cations are adjusted at pH 5.5 +lot
0
date. It was therefore necessary to test the electrode for its possible response to barium ion with acetone as the medium. The electrode was accordingly immersed in 90 vol YO acetone solutions with stepwise changes in their concentrations of barium chloride, and the electrode potential was measured each time after suitable equilibration. The results are shown in Figure 4 where measured potentials are plotted against pBa, the logarithm of the reciprocal concentrations of barium ion. For comparison with electrode response to monovalent cations, identical experiments were carried out with respect to hydrogen, sodium, and potassium ions with the results shown in Figure 4. The conclusion is inescapable that the electrode is not responsive to barium ion below 10-3M, even in a nonaqueous medium. Although the Nernstian slope was well reproduced for sodium and hydrogen ion in 90 vol ‘70 acetone, the order of selectivity for these two was reversed in a nonaqueous medium. The selectivity for hydrogen is now lower than that for sodium. This fact supports the hypothesis that hydrogen ion activity is considerably decreased in a high concentration of organic solvent (3). In view of the fact that the electrode did not show a Nernstian response to barium ion, a further study was made to account for the potential break in the sulfate titration. The electrode was immersed in a 90 vol ’70 acetone medium which was free of barium ion. After stirring well and letting the potential come to equilibrium, a drop of 0.005M barium chloride was added to the solution. In spite of the extremely small change in barium ion concentration (from zero to lO-eM), a characteristic response was observed. An electrode potential of about 10 mV instantly developed and then gradually subsided with time, approaching its initial level as shown in Figure 5 . Similar experiments were carried out in plain water and the results are shown in Figure 6 where curve I was obtained upon the addition of a drop of 0.005M barium chloride. Curve 11 was obtained by addition of two drops of 0.05M barium chloride with the second one being added 20 minutes after the first. Figure 6 shows that in water alone the larger change in barium ion concentration (from zero, or probably from any fixed level) evokes the larger change in potential, while the rate of change in concentration is effective in developing the potential in the same time. It must therefore be concluded that the potential break at the equivalence point in the sulfate titration is a transient response of the electrode to the barium ion. From a consideration of Figures 5 and 6, a simple interpretation of the titration curves in Figure 3 as far as their shape in 768
ANALYTICAL C H E M I S T R Y , VOL. 46, NO. 6, M A Y 1974
r
I
+ 2.3 x
~ O - ~ Bo2+ H
I 0
10
20
30
TIME, rnin
Transient responses of N A S 11-18 glass electrode to sudden change of barium ion concentration in aqueous medium at pH 6.5 Figure 6.
the vicinity of the end point is concerned, can now be attempted. In a highly nonaqueous medium like 90 vol % acetone, the barium ion concentration before the end point is extremely low on account of the low solubility product of barium sulfate, but as the end point is approached, the electrode begins to respond to the barium ion as soon as the rate of change in concentration becomes large enough to raise the potential. This gives rise to a gradual change in curve direction. But if the concentration of acetone in the medium is decreased slightly, e.g. to 70 vol %, the rate of change in the concentration of barium ion is smaller on account of the now slightly higher solubility of barium sulfate-ie., the number of Ba2+ ions being added is less significant in proportion to the number already present in solution. But when the equivalence point is finally reached, the sudden build-up of excess barium ions corresponds to a more significant rate of change and the transient response of the electrode appears suddenly. This produces a sharp change in curve direction. In a lower concentraqion of organic solvent-ie., 50 vol % acetone-the rate of change in barium ion concentration is still further diminished so that the potential jump becomes even less significant. The phenomena shown in Figures 5 and 6 seem analogous to the transient response reported by Rechnitz and his group who observed a development of potential for a sudden change in the concentration of calcium or strontium ion against a constant background concentration of potassium ion when using the Beckman “Cationic Glass Electrode.” The observed potential disappeared within a few milliseconds. They offered the following simple explanation for the mechanism of the transient response. A sudden change in the concentration of calcium ion in the
presence of potassium ion leads to ion exchange on the membrane surface of the electrode, and the difference in electrochemical potential between the two ions produces an instantaneous electrode potential. But the relatively low mobility of calcium ion in the hydrated glass layer of the electrode will not permit ion transport through the membrane so that the electrode potential immediately reverts to the initial value, which is merely a function of the potassium ion activity. One can therefore assume that a similar mechanism applies to the case of the hydrogen and barium ions when the monovalent cation-selective electrode is used in the titration of sulfate with barium ion. In the present case,
however, the transient response probably proceeds at a considerably slower rate than in the instance reported by Rechnitz. It may be expected that similar behavior by other divalent cations will doubtless be reported in the future. ACKNOWLEDGMENT The authors are indebted to J. A. Kuck for valuable discussions and helpful suggestions in the preparation of the manuscript. Received for review September 20, 1973. Accepted December 27,1973.
Polarographic Method for the Determination of Hexafluorosilicate lndira Rajagopalan and S. R. Rajagopalan Materials Science Division, National Aeronautical Laboratory, Bangalore 56001 7. India
Determination of hexafluorosilicate (SiF62- ) finds application in the control of baths for high speed chromium plating, black chromium plating, and winning of high purity chromium. In addition, fluoride is often determined as SiFe2- (after distillation) to eliminate interferences. The authors noticed that SiFs2- shifts the half-wave potential ( E l l 2 ) of the U(V) U(III) wave (supporting electrolyte, 0.5M KC1-15mM HC1) to more (-) negative potentials. This shift ( A E 1 / 2 ) varies linearly with the concentration of SiF62-. This relationship constitutes the basis of the present method.
-
EXPERIMENTAL Apparatus. A Metrohm water jacketed cell was employed. The temperature was maintained at 30 f 0.1 "C by circulating thermostated water through the water jacket. A saturated calomel electrode was used as reference electrode while the counter electrode was a mercury pool. A Sargent polarographic capillary was used. Normal and derivative DC polarograms were obtained with a Radelkis DC polarograph (Model OH-102). Drop times were mechanically controlled by employing a Radelkis Tastrapid adaptor (Type OH-991). This unit permitted recording of polarograms aided by sample hold circuitry. Reagents. Analytical grade reagents were used. Hexafluorosilicic acid (HzSiF6) and sodium fluoride were used as sources of SiF& and F- ions, respectively. Standard solution of HzSiF6 was prepared by distilling the acid and standardizing the distillate by potentiometric titration. Procedure. The solutions were prepared from twice distilled conductivity water. Deaeration was accomplished by bubbling for 30 minutes cylinder hydrogen gas purified by passing successively through activated carbon, vanadous sulfate solution and finally through 0.5M KC1. All the polarograms were recorded at a scan rate of 62.5 mV/min, with a drop time of 3 sec using a three-electrode configuration. From a derivative polarogram, El32 of the U(V) + U(II1) wave is obtained by measuring the peak potential (error, fO.l mV) from a recorded dc polarogram by the method of Hume, De Ford. and Cave ( I ) . In this work, the latter method was used. For the routine determination of (SiF&) in plating baths, the former method was employed. For estimating SiF& in chromium electrowinning and black chromium baths, where SiF+ is used as an addition agent, the following procedure has been used by the authors for two years. An aliquot suitable to give a SiF& (in the range of a ppm) in the final solution is transferred to a polarographic cell and 1 ml of
( 1 ) D. N . Hume, D. D . 73, 5323 (1951).
De Ford a n d C. E . Cave, J . Amer. Chem. SOC.,
hydrazine hydrochloride is added to reduce Cr(V1) to Cr(II1). Then amounts of KC1, HC1, and UO~(CzH30z)zare added such t h a t their concentrations in the final solution are, respectively, 0.5M, 15mM, and 0.5mM.
RESULTS AND DISCUSSION Kolthoff and Lingane (2) report that in moderately acid solution (0.01 to 0.2M HCl), the reduction of uranyl ion yields a doublet wave, whose half-wave potentials are -0.18 and -0.92 V us. SCE, independent of acidity, the concentration of uranyl ion, or whether KC1 is present in concentration up to about 0.5M. They identified the first wave to be due to the reduction of U(VI) to U(V) and the second wave to the reduction of U(V) to U(II1). Grether ( 3 ) gives a value of -1.25 V for E1 2 of the second wave. Bond and O'Donnell ( 4 ) report that E112 changes from this value a t very low acidity toward that given by Kolthoff and Lingane as acidity is increased. Our observations are similar to those of Bond and O'Donnell (Figure 1. Curve A ) . For our purposes, the absolute value of E1 2 is relatively unimportant, as we are only concerned with changes in Ellz. The present work has shown that addition of SiF& shifts E112 to more negative potential and decreases the id of both the waves of uranyl ion reduction-a result similar to that reported by Bond and O'Donnell on the effect of F- on the reduction waves of uranyl ion. The effect of U(II1) wave is more SiF+ on E1 2 and id of the U(V) marked than on the U(VI) U(V) wave. Further, it is noticed that changes in the value of Ell2 are more pronounced than the changes in id resulting from changing SiFs2-. Therefore, from an analytical standpoint. it will be better to employ Ell2 as a measure of SiFe2- and, hence, we will confine our discussions to the effect of SiFe2- on Ell2 of the second wave. It is well known that shift in Ell2 is associated with the formation of complex, inhibition or catalysis of electrode reaction ( 4 ) . The shift is normally in the negative direction in the case of complex formation ( 5 ) .Inhibition also
-
+
(2) I . M. Kolthoff and J. J. Lingane, "Polarography," Vol. 1 1 , 2nd ed., lnterscience, New York, N . Y . , 1952, p 462. (3) C. Grether, "Metrohm Application Bulletin," A36E, Jan 4, 1965. (4) A . M. Bond and O'Donnell, Ana/. Chem., 34, 1347 (1962). (5) J . Heyrovsky and J. Kuta, "Principles of Polarography," Publishing House of the Czechoslovak Academy of Sciences, Prague, 1965, p 148. A N A L Y T I C A L C H E M I S T R Y , V O L . 46, NO. 6, M A Y 1974
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