The order of elution of the silylated paromomycins I and I1 is analogous to silylated neomycin (7). It may be speculated, therefore, that the equatorial trimethylsilylamine group at the C 6position is exposed to the outer environment to a greater extent than the corresponding axial group of paromomycin I. As a result, paromomycin I1 is more soluble in the liquid support phase than paromomycin I. This coupled with the expected lower volatility of the equatorial form (9) would explain why paromomycin I1 is retained longer on the OV-1 column than paromomycin I. Separation of Kanamycins. Figure 5 shows a chromatogram of the kanamycins. The compounds were identified by the LKB 9000 gas chromatograph-mass spectrometer in order of elution as (1) aminoglucosyl deoxystreptamine, (2) kanamycin B, (3) kanamycin A. The fourth peak is assumed to be kanamycin C ; however, mass spectrometric analysis of this peak gave inconclusive data. The order of elution between kanamycins A and B is analogous to paromomycin I and neomycin B (Figure 4) and can be attributed to their structural differences which are at the C2 position of the glucosamine moiety (OH for kanamycin A us. NH2 for kanamycin B) (Figure 1). Characterization of Derivatives. Mass spectrometry was used to characterize silylated derivatives of paromomycin and kanamycin. Because silylated amines are readily hydrolyzed, drying and transfer of the derivatives into a mass spectrometer is extremely difficult. T o minimize hydrolysis, a sample was introduced into the chromatographic inlet system of an LKB 9000 gas chromatograph-mass spectrometer. Although the molecular ion of totally silylated paromomycin (1551) and kanamycin (1276) are beyond the optimum capability of the
LKB 9000, mass in excess of 1200 were obtained. The mass spectrum of paromomycin I exhibited strong intensities at 376, 449 or 450, and 726 mje indicating the presence of silylated deoxystreptamine, paromose and/or glucosamine, and paromobiosamine. These data indicate that all active hydrogens on both the hydroxy and amine groups in paromomycin are completely silylated. Active hydrogens on kanamycin were also silylated, since the mass spectrum of kanamycin A exhibited strong intensities at 376 and 451 mje indicating the presence of silylated deoxystreptamine and kanosamine. The mass spectrum of silylated paromamine gave the molecular ion of 899 mje. Quantitative Determination. PAROMOMYCIN. The precision of the gas chromatographic method was determined by assaying 8 separate preparations of paromomycin sulfate powder. The relative standard deviation of the quantitative determination of paromomycin I is 1.7% in a sample containing 10.7% paromomycin 11. The precision of the method was determined KANAMYCIN. by assaying 7 separate preparations of kanamycin sulfate powder. The relative standard deviation for the determination of kanamycin A is less than 1 %. ACKNOWLEDGMENT
The supply of paromomycin sulfate from Parke, Davis & Company and kanamycins A and B from Bristol Laboratories is acknowledged. E. J. Hessler is acknowledged for the supply of paromamine, paromomycins I and 11. P. B. Bowman and M. L. Knuth are acknowledged for the mass spectrometric analysis. RECEIVED for review June 18, 1970. Accepted August 24, 1970.
(9) W. J. A. Vander Heuvel, J. Chromatogr., 27,85 (1967).
Voltammetric Determination of Europium(ll1) Using the Lanthanum Hexaboride Electrode D. J. Curran and K . S . Fletcher IIP Department of Chemistry, Uniuersity of Massachusetts, Amherst, Mass. 01002
RECENTLY we described lanthanum hexaboride as an electrode for electrochemical studies( I ) and employed this material for the electrochemical generation of lanthanum(II1) for use in titrations ( 2 , 3) and as a polarized indicator electrode for acid-base titrations ( 4 ) . The present note describes the application of this electrode as a voltammetric indicating electrode for the determination of europium(II1). Few electrochemical studies of the europium system have been reported. Laitinen and Taebel (5) discuss a polarographic investigation using 0.1N NH4C1supporting electrolyte Present address, Research Center, The Foxboro Company, Foxboro, Mass. 02035. (1) (2) (3) (4) (5)
D. J. Curran and K. S. Fletcher 111, ANAL.CHEM., 40,78 (1968). Ibid., p 1809. Ibid., 41, 267 (1969). Ibid., 40, 1804 (1968). H. A. Laitinen and W. A. Taebel, ibid., 13,825 (1941).
and report the half-wave potential as -0.671 V us. SCE and the ratio of the average diffusion current to concentration, id/C, as 2.88 kA-l./mole. Using this value of id/C and the capillary characteristics given by these workers, the diffusion coefficient is calculated as 5.87 X 10-6 cm*/sec from the Ilkovic Equation. Gierst and Cornelissen (6, 7) also investigated this system polarographically using several supporting electrolytes. These workers report the half-wave potential to be -0.60 V us. SCE and the mean value (for several supporting electrolytes) of the diffusion coefficient to be 7.1 X 10-6 cm2/sec. Anderson and Macero (8) determined the formal potential of the system (Eo’= -0.60 V us. SCE in (6) L. Gierst and P. Cornelissen, Collect. Czech. Chem. Commun., 25, 3004 (1960).
(7) L. Gierst, in “Transactions of the Symposium on Electrode Processes,” E. Yeager, Ed., John Wiley and Sons, New York, 1961, pp 109-144. (8) L. B. Anderson and D. J. Macero, J. Phys. Chem., 1942 (1963).
ANALYTICAL CHEMISTRY, VOL. 42, NO. 13, NOVEMBER 1970
1663
Figure 1. Plots of Deak current US. concentration for reduction of europium(II1) ion at lanthanum hexaboride 0 Sweep rate = 8.90 mV/sec
V Sweep rate = 17.7 mV/sec Sweep rate = 44.0 mV/sec 0 Sweep rate = 88.5 mV/sec
0
0.2
0.4
0.6
0.8
1.0
1.2
1.4
1.6
1.8
2.0
CON C E N T RATION OF EdRI), mmoles/Li t e r
perchlorate media) using current reversal chronopotentiometry. Randles and Somerton (9) determined the apparent rate constant at mercury using the faradaic impedance method. No solid electrode voltammetric study of this system has been reported.
of the test solution. The electrode was then introduced into the cell, the cell was deaerated, and after return of quiet conditions, the electrode was switched from open circuit to -0.40 V us. SCE and the current-time or current-potential experiment was immediately begun.
EXPERIMENTAL
RESULTS
Reagents. Supporting electrolyte solutions were made from reagent grade KCI using water redistilled from alkaline permanganate and were pre-electrolyzed at a mercury pool held a t -1.25 V cs. SCE for several hours prior to use. Determinate europium(II1) solutions were prepared from Eu203 (material better than 99.9 EuZOB)which had been ignited to constant weight a t 900 "C in a platinum crucible. Weighed portions of the ignited material were dissolved with 25 ml of 6 N HCl and the solutions were evaporated to dryness to remove excess HCl. The europium chloride samples were then transferred to 100-ml volumetric flasks with weighed portions of recrystallized KCI. The p H of the solutions, as read with a standardized pH-glass calomel electrode pair and p H meter ranged from 5.1 to 5.5. Apparatus. The voltammetric cell, the electrodes, the potentiostat, and related circuitry have been described ( I ) . The lanthanum hexaboride rod had a geometric area of 0.317 cm2 and was used in its shielded configuration ( I ) . Currentpotential and current-time curves were recorded, respectively, o n a Moseley Model 135M X-YRecorder and a Heath Model. 20A Servo Recorder with a chart speed of 8 inches per minute. Sweep rate calibration was accomplished by recording the time for the sweep generator to sweep 1.00 volt, as noted on the calibrated X-Y Recorder, with a Model S-10 Precision Timer (Standard Electric, Springfield, Mass.). All potential calibrations were accomplished with a Model 2730 Portable Potentiometer (Minneapolis-Honeywell Regulator Company, Rubicon Instruments, Philadelphia). All work was done with the cell thermostated a t 25.00 f 0.02 "C using a P. M. Tamson (Zoetemer, Holland) Model T9 Constant Temperature Bath. The bath was turned off prior to each run to ensure quiet conditions. Prepurified nitrogen was used to free the cell solutions of oxygen. Electrode Pretreatment. The lanthanum hexaboride electrode was cleaned prior to each run by a brief washing with 6 N HC1 and then water, and finally was rinsed with a portion
Potentiostatic current-time curves were obtained for the reduction of 0.5229mM Eu(II1) in 0.100N KCI using the shielded lanthanum hexaboride electrode and quiet solutions a t 25.0 "C. For this work, the lanthanum hexaboride electrode was switched from -0.40 to -0.90 V cs. SCE and the current decay was recorded as a function of time for about two minutes. The currents were corrected for residual current, which was obtained using the same procedure but in the absence of Eu(II1). Values of it1'* were calculated for 1.50second intervals from t = 0 to t = 84.0 seconds and the average and mean deviation of each of four runs were 23.5 i 0.6, 21.6 + 1.0, 23.1 i 0.6, and 22.8 + 0.3 pA.sec*'?. The constancy of the product is taken as good evidence that the reduction of Eu(II1) at the lanthanum hexaboride electrode is diffusion controlled under these conditions. Evaluation of the diffusion coefficient of Eu(II1) from the linear diffusion equation using the average of the it112values shown above, n = 1 equiv;mole, F = 96, 487 coul,'equiv, A = 0.317 cm2, and C = 0.5229mM, gave D = 6.38 X 1 0 P cmL/sec with a relative deviation of 1 5 %. Current-potential curves were obtained for the reduction of 0.2814, 0.5354, 0.7701, 1.220, and 2.034mM E(II1) in 0.100N KCI using linear sweep rates of 8.90, 17.7, 44.0, and 88.5 mV/sec. In each case at least five curves were recorded. All curves were obtained using a n initial potential of -0.40 V us. SCE, and quiet solutions thermostated at 25.00 "C. Peak type current-potential response was obtained and the peak currents were corrected for residual currents by subtracting the residual current (obtained in prior sweeps) at the potential of the current peak from the peak current. Linear dependencies between peak current and concentration of Eu(II1) a t the four sweep rates were found and are shown in Figure 1. The averages and standard deviations of the i,'C ratios obtained at sweep rates of 8.90, 17.7, 44.0, and 88.5 mV/sec were 20 i 0.5, 28.6 =k 0.7, 46.6 1.5. and 66.7 + 4.8 PA-l./mole, respectively, showing an overall relative
(9) J. E. B. Randles and K . W. Somerton, Trans. Faraday SOC.,48, 937 (1952).
1664
ANALYTICAL CHEMISTRY, VOL. 42, NO. 13, NOVEMBER 1970
*
Table I. Voltammetric Potential Data for the Reduction of Eu(II1) at Lanthanum Hexaboride Sweep rate, mV/sec 8.9 17.7 44.0 88.5 - E p , V U S . SCE 0.665 0.671 0.685 0.700 1 0 . 0 0 3 1 0 . 0 0 3 1 0 . 0 0 2 10.004 0.602 0.605 - Ep,p, V U S . SCE 0.595 0.596 1 0 . 0 0 2 1 0 . 0 0 2 zt0.002 1 0 . 0 0 2 0.083 0.095 0.075 E p i ~- Ep, V 0.070 - Ep/o,8j17,V CS. SCE 0.627 0.630 0.638 0.646 10.002 10.002 10.002 10.002
100
-
90 -
,g
80-
4
70-
0 I '-I
z 60W
Table 11. Ratios of Experimental Peak Currents to Peak Currents Calculated for a Reversible Charge Transfer Sweep rate, mV/sec 8.90 17.7 44.0 88.5 C = 0.2814mM/I, (ip)exptd(ip)oaicd 0.953 0.988 1.072 1.149 C = 0.5354mM/l, (ip)expt~~/(ip)eaied0.972 0.999 1.051 1.091 C = 0.7701mM/1, 1.050 1.053 1.021 1.040 (ip),.,t~i/(ip)oaiod C = 1.220mM/I, 0.998 0.959 1.002 (i,)ex,t~i/(ip)oaicd 1.006 C = 2.034mM/1, 0.996 0.959 0.985 0.974 (ip)expt'l/(ip)calcd
precision of &4%. Linear dependencies between peak current and square root of sweep rate were found and are shown in Figure 2. The averages and standard deviations of the ip/V112ratios obtained using the 0.2814, 0.5354, 0.7701, 1.220, and 2.034mM Eu(II1) solutions were 62.9 1 5.2, 118 i 7, 172 i 3, 260 i 5, and 483 i 3 pA.sec1/*/volt1/*, respectively, showing a n overall relative precision of 3 %. Experimental values for the peak potential, E,, the halfpeak potential, Epll, and the potential a t which the current is 85.17% of the peak current, E p / ~ . 8 5 1did 7 , not show dependency o n concentration of Eu(II1) but did show cathodic shifts with increasing sweep rates. The values, together with EPi2- E, values, are shown in Table I. I t has not been possible t o arrive a t kinetic parameters for the europium reduction a t LaBGo n the basis of these data because the mechanism remains unknown. On the assumption that it might be a simple quasi-reversible charge transfer case as has been reported o n mercury (6, 7, 9), a curve fitting approach was adopted using the computer program of Nicholson for the corresponding cyclic case but omitting the reverse scan in the calculations (IO). This approach is necessary because Nicholson has shown that the curve shape depends o n both the transfer coefficient and the electron transfer rate constant. F o r the calculations, the ratio of diffusion coefficients was taken as 1.155 on the basis of the work of Gierst and Cornelissen (6). No combination of QI and k , was found to account for the experimental data. The best fit obtained yielded Epl? - Ep separations which agreed with the experimental values within several millivolts but the calculated peak currents were always 10 to 20% too low. Further, plots of i p / V 1 / 2 us. V and AE,lAlog V us. V did not produce a diagnosis consistent with that for any mechanistic scheme characterized by Nicholson and Shain (11). Table I1 shows the ratios of experimental peak currents t o theoretical peak currents for a simple reversible electron transfer process. The trend is an increasing ratio with increasing sweep rate for the lower three concentrations, but appears to
a
K
2
50-
0
0.10 0.20 SQUARE R O O T O F SWEEP RATE, (VISEC.)'"
Figure 2. Plots of peak current VS. square root of sweep rate for reduction of europium(II1) ion at lanthanum hexaboride V 0.2814mM Eu(II1) 0 0.5354mM Eu(II1) V 0.7701mM Eu(II1)
*
(10) R. S. Nicholson, ANAL.CHEM., 37,1351 (1965). (1 1) R. S. Nicholson and I. Shain, ibid., 36,706 (1964).
0.30
0 0
1.220mM Eu(II1) 2.034mM Eu(II1)
be a decreasing ratio for the remaining concentrations. F o r the two highest sweep rates, the trend is to lower ratios with increasing concentration while for the remaining two sweep rates, the ratios go through a maximum with increasing concentration. It is clear that the mechanism cannot be described by either a reversible or a quasi-reversible charge transfer process. It seems more likely that the mechanism is complex. Several facts support this contention. It was observed o n repeated scans without any electrode pretreatment between runs that the peak current behavior would disappear yielding to a limiting current behavior which, in turn, would decrease o n further scans and nearly disappear. Evidently, a filming process is in operation. It has been observed in these laboratories that a similar film is formed, along with gas evolution, a t a large carbon cathode during the constant current electrolysis of Eu(II1) in a perchlorate medium a t p H 5 t o 6 under conditions where 100% current efficiency is expected (12). It is believed that the film is E U ( O H ) ~ . The reaction, Eu(I1) 2H+ -,Eu(II1) HY, is well known and believed to be fairly slow in this p H region. However, Biedermann and Terjosin report that the reaction is catalyzed by platinum ( 1 3 ) . We suggest that the reaction is also catalyzed by carbon and by LaB8. Thus a plausible explanation for the facts is the following: the charge transfer process is
+
+
(12) D. J. Curran and L. Jaycox, unpublished data, these laboratories. (13) G. Biedermann and G. S . Terjosin, Acta Chem. S c a d , 23, 1896(1969). ANALYTICAL CHEMISTRY, VOL. 42,
NO. 13, NOVEMBER 1970
1665
transiting from a reversible toward a totally irreversible process in the range of sweep rates investigated here but the Eu(I1I) is being regenerated chemically. This would account for the shift in potential and the higher peak currents. However, the chemical process may eventually make the solution basic enough in the vicinity of the electrode surface to precipitate Eu(II1) as the hydroxide. This would explain the filming and the decrease in peak currents. The relative rates of the chemical reactions and their speed with relation to the sweep rate would determine the nature of the current-potential curve. For the present, the linear dependence of peak current o n concentration and o n the square root of the sweep rate as
shown in Figures 1 and 2 must be viewed as empirical but useful data.
ACKNOWLEDGMENT We thank J. E. Roberts of this department for providing the europium oxide from his private stock and also F. O’Brien, R. L. Myers, E. Rickard, and L. Jaycox for their assistance with the computer program. RECEIVED for review January 22, 1970. Accepted August 24, 1970. Presented in part at the First Annual Meeting, Northeastern Section of the A.C.S., Boston, Mass., October 1968. Taken in part from the Ph.D. thesis of K.S.F. 111.
Rapid Desorption of Chromium(ll1) from Cation Exchanger with Hydrogen Peroxide Solutions Gerald A. Sleater and David H. Freeman National Bureau of Standards, Washington, D . C.20234
INTHE PRESENT STUDY we are concerned with the removal of Cr(II1) from a cation exchanger in order to obtain the release of precisely known amounts of chromium from the corresponding ion exchange particles. The individual particles are spherical, and therefore microscopically measurable, resin beads of homogeneous matrix, which can be used as calibration microstandards ( I , 2). When the exchanger is so used, it is important that desorption be accomplished with minimum contamination from the desorbing agent. F o r multivalent, strongly sorbed ions, simple electrolytic reactions may be kinetically limited and more rigorous techniques become necessary. The thought of converting cationic Cr(II1) into anionic Cr(V1) suggested the possibility of using a n oxidative desorption process. Hydrogen peroxide, with its available high purity and innocuous oxidation products, was a n obvious choice. Preliminary qualitative experiments showed that a basic solution of hydrogen peroxide would remove the sorbed Cr(III), the yellow color of Cr04*- being immediately visible. The use of hydrogen peroxide in elution of ion exchange resins has been limited to either stabilizing an oxidation state in solution (Nb”, Tav) (3) or acting as a complexing agent (Ti) ( 4 ) . Its previous use in the oxidation of a sorbed-species in situ has been reported only for Fe(I1) being oxidized to Fe(II1) which was then desorbed (5). In the following series of batch experiments, hydrogen peroxide solutions at various pH’s were reacted with a Cr(II1) perchlorate loaded cationic exchanger. The completeness of Cr removed has been evaluated under varied conditions. The methods were not tried with packed columns because of (1) D. H. Freeman and R. A. Paulson, Nature, 218, 563 (1969). (2) D. H. Freeman, L. A. Currie, E. C. Kuehner, H. D. Dixon, and R. A. Paulson, ANAL.CHEM., 42, 203 (1970). (3) F. W. E. Strelow and C . J. C. Bothma, ibid., 39, 595 (1967). (4) D. I. Ryabchikov and V. E. Bukhtiarov, 2. Analit. Khim., 15 242 (1960); C . A . , 54, 19261f(1960). ( 5 ) A. Rusi and S . Ionescu, Acad. Repub. Pop. Rom. Stud. Cercer. Fiz., 14, 127-132 (1963); C.A., 59, 5754e (1963). 1666
the gas evolution that accompanies the side reaction of H20? decomposition.
EXPERIMENTAL Reagents. Reagents were prepared from ACS Reagent Grade material. The Cr(II1)per chlorate solution used in loading the cation exchanger was prepared starting with chromium sesquioxide. All hydrogen peroxide solutions were made from the same lot of 50x hydrogen peroxide (stabilized), which was stored at -25 “C to retard decomposition. Analysis. p H was measured with a commercial p H meter and glass electrode-saturated calomel electrode system, checked against a standard buffer solution. In all analyses, chromium was determined spectrophotometrically as chromate in basic solution (6). Measurements were made at room temperature (21-24 “C) with recording spectrophotometers using matched quartz cells. Extinction coefficients were measured at 375 nm in the visible spectrum and compared with a calibration curve prepared for the range 1-10 ppm Cr. Blanks were run on all reagents used and indicated that no correlation factors were necessary within the limits of the spectrophotometric methods used. Resin. Fifteen grams of a dry, hydrogen-form sulfonic acid exchanger (8% crosslinking, 5-25 pm size) was column loaded with an acidic (pH 1.8) 0.05MCr(III) perchlorate solution. The converted resin was agitated and washed with distilled water, washed with absolute ethanol to shrink the exchanger, and dried overnight a t 85 “C under an oil pump vacuum of 50 torr. In a quantitatively loaded exchanger, the weight fraction of chromium (f&) can be calculated from the equivalent weights of the exchanger network (Eo)and the counterion ( WC,);the equivalent weight of the exchanger had been previously determined to be 198 g/equiv (7). From the formula
+
fcr = ECr/(Eo Ec~) the weight fraction of Cr(II1) in the exchanger is predicted to be 0.0804 (8.04 wt %).
(6) G. Charlot, “Colorimetric Determination of the Elements,” Elsevier, New York, N. Y . , 1964, p 227. (7) See Ref. 2.
ANALYTICAL CHEMISTRY, VOL. 42, NO. 13, NOVEMBER 1970
