884
ANALYTICAL CHEMISTRY
acetone was added, and the mixture was magnetically stirred for 5 minutes. When the emulsion settled, the acetylacetone floating on top was carefully siphoned off. This extraction procedure was repeated seven times. After the fifth extraction the blue color of cop er was not evident in either phase. On completion of the seventt extraction, the aqueous phase was extracted with carbon tetrachloride to remove excess acetylacetone and then made ammoniacal. No blue color due to the cupric ammonium complex was seen. The zinc remaining in the aqueous layer was determined with 8-quinolinol. Weight of zinc recovered was 0.1315 gram. Continuous Extraction of Copper. Twenty-five milliliters of a solution containing 0.1190 gram of zinc ion, 25 ml. of a solution containing 0.100 gram of copper ion, and 250 ml. of distilled water were added to the extraction bulb of an apparatus designed for continuous extraction with solvents lighter than water. The solution contained sufficient sulfuric acid so that when diluted to 300 ml. the pH was 2.4. The pH was determined on a separate identical solution. Two hundred and fifty milliliters of acetylacetone was placed in the solvent boiling flask, and the extraction was begun and continued for 4 hours. The aqueous phase became colorless in 2.5 hours. At the end of the extraction process the pH was measured and found to be 1.9. The aqueous phase was poured into a 500-ml. volumetric flask together with the water used for washing the bulb, and the solution brought to volume. Exactly 250 ml. of the solution was reserved for the determination of the zinc, and the remainder was used for a test of any remaining copper. Test for Copper. The 250-ml. remainder of the aqueous phase was acidified with concentrated hydrochloric acid until the pH was 1.0. It was shaken with four successive portions of carbon tetrachloride to remove the excess acetylacetone. It was then shaken with 10 ml. of a dilute solution of dithizone in carbon tetrachloride so prepared that 20 micrograms of copper would change the color from green to red. Less than 20 micrograms of copper were found; therefore the aqueous phase contained less than 40 micrograms of copper i n entirety. Determination of Zinc. Because the presence of a large excess of the acetylacetone was found to cause slightly low results in the determination of zinc by 8-quinolino1, the 250 ml. of the aqueous phase reserved for the determination of zinc was boiled for 1 hour to remove the excess acetylacetone. Zinc was determined by the
8-quinolinol method described above. The total amount of zinc remaining in the aqueous phase was found to be 0.1183 gram as compared with 0.1190 gram originally present. These values probably agree within the experimental errors inherent in the 8quinolinol method. DISCUSSIOY
By extraction with acetylacetone, copper can be quantitatively removed from a zinc-copper mixture without removing any of the zinc. I n itself this separation may not be of great analytical significance; however, as continuous extraction with acetylacetone is readily accomplished when the pH,/$ values are well separated, and over 60 metals form chelates with acetylacetone, the method xhen extended might well serve as the basis for a systematic separation scheme for the metals. ACKNOWLEDGMFJVT
The authors gratefully acknowledge the financial support of the U. S. Atomic Energy Commission. LITERATURE CITED
(1) Furman, K. H., Rlason, W. B., and Pekola, J. S., AXAL.CHEM.,
21, 1325 (1949). (2) Irving, H. M., and Williams, R. J. P., J . Chem. Soc., 1949,1841. (3) Kolthoff, I. lI,,and Sandell, E. B., J. Am. Chem. Soe., 63, 1906
(1941). (4) AIoeller, T., ANAL.CHEM.,15, 346 (1943). (5) S a c h o d , F. C., 2. p h y s i k . Chem., A182, 193 (1938).
(6) Rydberg, J., Ssensk Kern. Tidsk., 62, 179 (1950). (7) Sandell, E. B., “Colorimetric Determination of Traces of Metals,” p . 300, New York, Interscience Publishers, 1950. (8) Van Uitert, L. G., Fernelius, W. C., and Douglas, B. E., C . S. Atomic Energy Commission, R e p t . NYO-3370. RECEIVED for review October 13, 1952. Accepted February 24, 1953. Presented a t the h l e e t i n g - i n - ~ ~ i n i a t uof r e the Pittsburgh Section, CHE\IIC.AL SOCIETY, Pittsburgh, Pa., June 6, 1932.
.4lIERICAN
Voltammetric Determination of lead as lead Dioxide at the Rotated Platinum Wire Electrode I. 31. ICOLTHOFF, JOSEPH JOKDAR”,
AND
ALBIN HEYNDKICKX
School of Chemistry, rnicersity of Minnesota, Minneapolis, Minn.
I
S ALKALI hydroxide solutions thallous thallium yield6 anodic
waves a t the rotated platinum electrode ( 1 ) with well defined diffusion currents ( 3 ) . The Tyave height corresponds to the 2electron transfer thallium(1) to thallium(III), the diffusion current being proportional to concentration. I n the present study the osidation of lead(I1) to lead(1V) has been investigated at the rotated platinum electrode and suitable conditions have been found for the determination of lead in very dilute solutions by measuring its anodic diffusion current. EXPERI MEYTAL
chemicals and conductivity water were used throughout. A standard solution of 0.005 X lead acetate was made up in thin 0.01 .lf perchloric acid. This stock solution was used to prepare concentration range. lead solutions in the IO* M to 10-3 Current-voltage curves were recorded automatically with a Sargent Model XXI polarograph, the rate of voltage change, AE/At, being 3.70 mv. per second. Currents were sometimes measured manually at predetermined potentials using a circuit previously described ( 3 ) . d simple electrolysis cell (3)made of a 100-ml. borosilicate glass beaker was used, connected to a saturated calomel electrode by means of a Hume-Harris salt bridge ( 2 ) . C.P.
1
O n leave from the Hebrew University, Jen,-aleni, Israel
Indicator electrodes were made of platinum wire, 0.5 mm. in diameter and approximately 5 mm. in length, and rotated a t 600 r.p.m. Most of the results reported in this paper were obtained with an electrode which yielded for the thallous-thallic oxidation in 0.1 M sodium hydroxide an i d / c value of 218 iz 2 pa. per millimole per liter a t 25’ C. In some experiments electrodes of slightly different size were used. The values of all diffusion currents were referred to the above electrode by making use of the fact that diffusion currents corresponding to a given electrode process measured a t different electrodes a t the same rate of stirring are proportional to the surface area of the electrode. Thus the ratios of diffusion current constants, i d / c , measured a t three different electrodes were found to be the same a t each electrode for the following reactions: thallous to thallic, thallous to thallium, Ag(NH3)2+ to silver, and oxygen to hydrogen peroxide. The platinum wires were sealed into glass, and electrical contact Tvas made with the aid of mercury. The platinum-glass joints were carefully annealed, because irregularities a t the metalglass interface give rise to abnormally large residual currents. Reproducibility of limiting currents depended greatly on the cleaning and pretreatment of the electrodes. The following procedure was found to give satisfactory resulta.
V O L U M E 2 5 , N O . 6, J U N E 1 9 5 3 On the basis of previous studies of the thallousthallic system it was anticipated that the anodic oxidation of divalent lead at the rotated platinum electrode could be used for rapid voltammetric determinations. Anodic current voltage curves of divalent lead have been studied in various media with an automatic recording polarograph. Well defined waves are found in dilute alkali hydroxide solutions. The electrode reaction is given by OH- -S PbOs(S) Hz0 2e. The anodic HPb02diffusion currents are proportional to the concentration of lead in the range betw-een and M. The standard potential of the lead dioxide-plumbite system has been calculated from voltammetric data. The lead(I1)-lead dioxide waves are ideally suited for the determination of lead in very dilute solutions. In 0.05 M sodium hydroxide as supporting electrolyte lead is determined with an accuracy and preM solution. The diffusion curcision of 29'0 in rent can be measured manually at a potential of $0.65 volt us. SCE. A procedure is described by which the determination can be carried out in the presence of at least a hundred fold excess of thallium.
+
+
+
The electrode was washed in 10% hydrogen peroxide plus 1 Af hydrochloric acid (to dissolve lead dioxide) and subsequently in 5 hf nitric acid. In this way very small residual currents were always observed within the range of potentials used for the determination of lead (see figures). The reported diffusion currents were found by subtracting residual currents from total limiting currents. All experiments were carried out a t 25.00' j=0.02" C . Potential values reported in this paper are referred to the saturated calomel electrode, unless otherwke stated. RESULTS
*inodic waves of lead with well defined diffusion currents were obtained only in dilute (0.01 to 0.1 M) alkali hydroxide solutions. In perchloric acid and in acetate buffers of pH 4 the residual current curves virtually overlap those of 10-4 A f lead solution.
885
Alkali carbonates and borates are not suitable media because of the limited solubility of lead in solutions of these electrolytes. Addition of a complex former, such as tartrate, increases the solubility of lead. However, diffusion currents of lead in borax and tartrate become less well defined with increasing concentration of tartrate, the anodic wave being drawn out toward more positive potentials and hardly separated from the decomposition current of the supporting electrolyte. h wave with a very narrow diffusion current region is obtained, for instance, in 0.05 Ji' borax plus 0.005 M tartrate. It was anticipated that well defined waves might be obtained in ethylenediaminetetraacetate media of varying pH, as quadrivalent lead is expected to form a more stable complex than divalent lead (IO). Indeed, a current-roltage curve of 10-4 Ji lead in 0.01 M sodium ethylenedianiinetetraacetate solution of p H 5.3 exhibits an anodic vave with a limiting current extending over the potential range between 0.7 and 1.1 volt. However, the corresponding residual currents are so large as t o cause appreciable uncertainties in the values of the diffusion currents. In ethylenediaminetetraacetate solutions of pH 7 and 10 the residual current lines and the current-voltage cur\-e of 10-4 Jf lead nearly coincide. In the chelating media (tartrate and ethylenediaminetetraacetatej no deposit was formed on the platinum wire anode, the lead( 11') formed remaining in solution. From solutions of lead in alkali hydroxide, a yellowish-brown coating is deposited on the electrode. Current-voltage curves of 10-4 31 lead in dilute alkali hydroxide solutions are plotted in Figure 1. A comparison of the current-voltage curve in 1 -1f sodium hydroxide (curve I) and the corresponding residual current (curve 1') shows that lead is heing oxidized a t potentials more positive than +0.2 volt. Thi- is further substantiated by the concomitant appearance of a broa n deposit on the electrode. However, no diffusion current region is obtained. In 0.1 J1, 0.05 -11,and 0.01 hf sodium hydroxide (Figure 1, curves 11, 111, and IVj well defined anodic waveE are obtained with practically equal diffusion currents. Current voltage curves obtained a t a lead dioxide-coated electrode, started a t +0.2 volt, mere identical uith the corresponding portion of the waves obtained a t the initially blank platinum electrode.
Table I.
Half-Wave Potential of Anodic Lead Wave i n Sodium Hydroxide CNaOH
POTENTIAL
,VOLT
vs
SC E
Figure 1. Current-Voltage Curves of 10-4 41 Lead in Sodium Hydroxide
CPb(1I)
El, z (SCE)
The waves are shifted but slightly toward positive potentials n-ith decreasing concentrations of the alkali, while the oxygen evolution potentials of the supporting elertrolytes are shifted in the same direction, but to a sonienhat greater extent. Conrequently, the diffusion current extends over the widest potential range in 0.01 M sodium hydroxide (from +0.6 to f0.8 volt), while in 0.1 ,?.I alkali the range narrows to between +0.55 and +0.7 volt. Plots of concentration of lead L I S . diffusion current (measured a t a potential of +0.65 volt) in 0.1 M, 0.05 Jf, and 0.01 M alkali, respectively, are presented in Figure 2. As can be seen from the figure, proportionality between id and CPI, is limited in 0.01 -11 sodium hydroxide up to a concentration of about 2 X rM lead. In 0.05 -Vi' and 0.1 M alkali diffusion currents and concentrations are proportional in the range investigated-Le., between 10-6 M and l O - 3 M . These two supportiiig electrolytes are recommended for analytical use. Half-n a r e potentials a t varying concentrations of lead and alkali are listed in Table I. These values are not measurably affected by addition of sodium perchlorate to keep the ionic strength in all solutions equal to 0.1.
ANALYTICAL CHEMISTRY
886
All current-voltage curves described above Eere run from zero toward positive potentials. When run in the opposite direction, from S0.65 volt to zero, a dissolution pattern (3, 9) such as the one shown in Figure 3 is obtained. If the carrent-voltage is started a t a more positive potential, where oxygen is evolved from the supporting electrolyte, and is run to potentials as negative as -0.3 volt a second dissolution pattern is obtained a t this negative potential. The second pattern is small in area and is attributed to the reduction of oxygen sorbed on the electrode surface. In this experiment dissolved oxygen was removed from the solution by deaeration with purified Linde nitrogen.
Dissolution Pattern of Lead Dioxide. The location of this pattern may serve for the detection of extremely thin films of lead dioxide, while the recorded current-time area provides a quantitative measure of the amount of dioxide originally present in the film ( 4 9 ) . ELECTRODE REACTION
The ionic state of divalent lead a t hydroxyl ion concentrations from 0.01 to 1 h ' corresponds practically quantitatively to t h e biplumbite ion, as has been shown by Lingane (8) and is borne out by the relevant thermodynamic equilibrium constants (6). Consequently, the anodic oxidation of lead a t the rotated platinum electrode may be formulated as:
150
HPbO;
100
HPbO;
a
L!
E,.. . = const.
Y
"3
+ 2e + H2O -+- Pb + 3 0 H -
(2)
A wave corresponding to the reversible process (Equation 1) should be determined by an equation of the form:
+
f
(I)
This accounts for the equality of the anodic and catbodic wave heights, the corresponding cathode process also involving two electrons (8):
W VI
a p"
+ OH- +PbOl (S) + H20 + 2e
50
0
10'
OXIO"
3XIO"
MIO.'
CONCENTRATION
5 X d
6XlO4
7x10''
@%IO.'
0x10.'
lo-'
OF LEAD, MOLES / L I T E R
Figure 2. Plot of Anodic Diffusion Current us. Lead Concentration in Alkali Hydroxide Solutions
A current-voltage curve of 10-4 , I Ilead in oxygen-free 0.05 M Bodium hydroxide solution run from zero to negative potentials yielded under experimental conditions a cathodic wave [lead(11) to Pb'] of height (13.8 pa.) equal to that of the anodic lead wave. The half-wave potential of the cathodic wave was about -0.8 volt. APPLICATIONS
Voltammetric Determination of Lead. From results plotted in Figure 2 it appears that the value of i d / C is the same in 0.1 i1f a~ in 0.05 Ai sodium hydroxide and equal to 138 pa. per millimole per liter, the standard deviation being A3 pa. in 0.1 Af and *2pa. in 0.05 M alkali. For quantitative purposes 0.05 JI sodium hydroxide is recommended as the supporting electrolyte. Procedure. In a given volume of the sample, which should not be less than 5 ml., adjust the lead concentration to a value between about 10-6 and 10-3 M. Add sodium hydroxide to make solution about 0.05 M in alkali. Measure the limiting current a t + O . G volt and correct for the residual current. Do not electrolj-ze the solution longer than 1 minute, in order to keep depletion of lead negligibly small. Oxygen need not be removed. Applying this procedure to an unknown solution containing 113 mg. of lead per liter, correct values were found with &2%. Determination of Lead in Presence of Thallium. Thallous ion is oxidized in alkaline solutions a t the rotated platinum electrode, yielding an anodic wave which interferes with the determination of lead (3). To determine lead in the presence of thallium the following procedure was developed. Acidify the solution to 0.1 hi in perchloric acid; add chlorine water to oxidize thallous thallium to thallic; remove excess chlorine by bubbling with nitrogen; add sodium hydroxide to make solution 0.05 M in alkali; and determine lead as described above. The thallic oxide, which precipitates when the alkali is added does not interfere and need not be removed. Using this procedure, lead in a concentration of lo-' M was determined in the presence of 0.01 di thallium with the same precision and accuracy as in the absence of thallium.
- 0.030 IOg ( i d - i)
(3)
where E,.*. denotes the potential of the indicator electrode. Plots of E, e. us. log (id - i) for the three anodic waves shown in Figure 1 yielded straight lines but with varying slopes. This indicates that the currents on these waves (and on the similar waves obtained a t an electrode coated with lead dioxide) may be determined both by the rate of diffusion and by the rate of the electro-ouidation. If the wave u ere reversible, a tenfold change in the activity of either the hydroxyl or of biplumbite ion should cause a shift of half-wave potential of +0.030 volt. It is seen from Table I that the observed half-wave potential shift is in the predicted direction but, as expected, the changes do not correspond quantitatively to those calculated. The relatively slow rate of electro-oxidation would also account for the fact that no defined wave is obtained in 1 ill sodium hydroxide. The i d / C values for the thallous-thallic and the plumbite-lead dioxide ovidations in 0.1 M sodium hydroxide were 218 and 138, respectively, corresponding to a ratio of 1.58. At 25" C. the difand that of fusion coefficient of the thallous ion is 2.00 X or a ratio of the biplumbite ion 0.98 X lO-5sq. cm. X see.-' (4,5) 2.01. The power of this ratio is 1.50. Thus the ratio of the cathodic and anodic diffusion .4O currents of thallium L.l and lead a t the +a0 w r o t a t ed platinum electrode varies p .eo under experimental 0 conditions with +- . I O z D2I3[cf. ( 7 , II)]. Y a I n a previous a 0 paper (3) it was shown that a sys-10 tem can be reversible and in equilib-20 rium with the I I I I I electrode a t t h e .IO '08 +om +a* .oz 00 POTENTIAL, VOLT vs SCE "zero c u r r e n t potential"-ie., at Figure 3. Dissolution Pattern of t h e potential a t Anodic Deposit which a dissolution Current-voltage curves of 10-4Mlead in 0.1 M sodium hydroxide run from $0.65 t o 0 pattern curve intervolt sects the residual Dotted line indicatea residual current
s
V O L U M E 25, NO. 6, J U N E 1 9 5 3 Table 11. csaon
conditions, the average value of E” being 0.22 us. NHE volt comp u e d to the calculated value of 0.21 volt.
Oxidation Potential of Lead Dioxide as Function of C N ~ O and H CPb(I1) CPb(I1)
0.1 0.1 0.01 0.01
a87
10-5
10-4 10-6
lo-‘
Zero Current Potential (SCE) +O. 160 0.12 0.19; 0. 155
ACKNOWLEDGMENT Eo(us. NHE)O
Acknowledgment is made to the Graduate School of the University of Minnesota. for a grant in support of this work.
$0.221 0.216 0.220 0.220 0 . 2 1 9 =t0 . 0 0 2
Mean Using Eilriation 4 and artivity coefficients of hydroxyl and biplurnhite on.$ of 0.81 and 0.89 in 0.1 M a n d 0.01 I\. no4ium hy,iroxide, respectively. a
LITER.ATL‘RE CITED
Delahay, P., and Stiehl, G. L., J . Am. Chem. Sac., 73, 1755 (1951 ).
Hume, D. iX , and Harris, W.E., ISD.ENQ.CHEM.,AXAL.ED., 15,465 (1943).
current curve of the supporting electrolyte-even if the corresponding waves are partly rate-controlled. If this were true for thepresent system, the zero current potential (or oxidation potential), E,=o,of the lead dioxide electrode should vary according to the rehtion :
E,,o
=
E”
- 0.030 log U O H -
- 0.030 log
UHP~OP-
(4)
E” in Equation 4 should be identical with the standard potential of Reaction 1, which can be calculated accurately from free energy data. Vsing Latiiner’s (6) values with the polarographic sign convention, we calculttte E” = +0.207 volt os. normal hydrogen electrode. I n Table I1 are given the authors’ data of potentids of lead dioxide a t varying alkali and lead concentrations. It is seen that Equation 4 accounts satisfactorily for the potentid of the plumbite-lead dioxide system, under experimental
Kolthoff, I. M., and Jordan, J., J . Am. Chem. Sac., 74, 382 (1952).
Kolthoff, I. hI., and Lingane, J. J., “Polarography,” 2nd ed , Vol. I, pp. 52,217, Kew York, Interscience Publishers, 1952. Ibid., Vol. 11,p. 528.
Latimer. W.bi.. “Oxidation States of the Elements and Their Potentials in Aqueous Solutions,” 2nd ed., pp. I50 ff., Kew York, Prentice-Hall, Inc., 1952. Levich, B., Acta Physicochim. U.R.S.S., 17, 257 (1942); 19, 117 (1943): Discussions Faraday Soc.. 1 , 3 7 (1947).
Lingane, J. J., Chem. Revs.. 29, 1 (1941). Lord, S. S., Jr., O’iXeill, R. C., and Rogers, L. B., ANAL.CHEY., 24,209 (1952).
Schwarzenbach, G., and Freitag, E., H F ~Chirn. . Acta, 34, 1492 (1951).
Tsukamoto, T., Kambara. T., and Tachi, I., Proceedings of 1st International Polarographic Congress, Prague, Vo1. I, p. 525,1951. RECEIVED for review November 24, 1952. Accepted January 1 6 , 1953.
Suppression of Cyanogen Bands in the Direct Current Graphite Arc by lithium Chloride ROBERT G. KEENAN’ AND CHARLES E. WHITE University of Maryland, College Park, M d . The production of extremely dense cyanogen bands in the 3500 to 4800 A. region of the emission spectrum, when a graphite arc serves as the excitation source, has rendered many valuable persistent line spectra valueless for qualitative and quantitative spectrographic analysis. lllethods to circumvent this difficulty include employment of less sensitive lines in other regions of the spectrum, use of the generally less sensitive spark source, use of metallic electrodes, at a great loss of sensitivity, or exclusion of nitrogen by maintenance of an atmosphere of steam, oxygen, or helium around the arc. None of these methods is completely satisfactory for routine
A
31 IJOR difficulty, encountered in the use of graphite electrodes in spectrochemical analysis, is the production of cyanogen band spectra, when graphite is burned in air, due to the formation and subsequent excitation of the cyanogen radical in the atmosphere of the arc. These bands, occurring in the 3500 to 4800 -4.region, are five in number and emanate from their respective “band heads” a t 3590, 3883, 4216, 4600, and 4740 A. The last two are produced only under severe conditions of expoNure. The first three are the most prominent, the most extensive, and the most readily produced. They are very dense and their fine line structure extends throughout the 3500 to 4216 region. Theqe bands mask many valuable persistent line spectra, thereby making difficult the qualitative identification of minor and trace quantities of metallic elements in this region. As most of this 1 Present address, Division of Occupational Health, U. S. Public Health Service. Cincinnati, Ohio,
determination of trace elements. The method described employs a lithium chloridegraphite matrix and permits qualitative and quantitative spectrography throughout the cyanogen band region. It is suitable for routine application. Cyanogen bands have been suppressed sufficiently to permit quantitative spectrography of vanadium, molybdenum, and titanium, employing lines which occur normally within and between band structures. Triplicate analyses show that good results are obtained over the electrode content range of 0.1 to 3.0 micrograms. The function of the lithium chloride is to lower the potential drop across the arc.
portion of the spectrum iu well blackened by these bands, quantitative analysis by means of lines in this region is practically impossible when conventional procedures are employed. Elements whose most persistent lines occur within these band structures are determined quantitatively in most cases by the use of less sensitive lines of the element in other portions of its spectrum; the employment of the spark source rather than the arc, with the production of less dense cyanogen bands but at a sacrifice of sensitivity; or the use of metallic electrodes, such as copper or silver, which also reduce the intensity of cyanogen band spectra, but a t a great loss of senqitivity. I n order to utilize this spectral region for qualitative and quantitative analysis, two previous methods have been developed for the suppression or elimination of cyanogen band spectra. T h e first is that of Ashton ( I ) , who found that lead oxide, when mixed with plant ash samples, upon exposure in a 6-ampere. 4,s-
