metric technique were in closer agreement with the quantities added than those by the spectrographic method. Thus, this detection method appears applicable to a practical analytical method for rubidium. However, any practical flame photometric method for the determination of rubidium should include an internal standard (such as lithium) and excess potassium on samples that contain variable potassium lwels. In this work, the standard curves were obtained from calibration solutions containing excess potassium only. Over a \Tide range of high potassium levels, Li R b ratios remain constant (Tables I1 and 111); in addition, it appears that thc excess potassium present acts as a predominant ion and minimizes
variations caused by other ions. Consequently, the use of calibration standards containing only one high level of potassium seems justified, provided the samples also contain a large quantity of potassium (above 20 mg. of potassium per ml.). Although this flame photometric study is based on solutions of known mineral composition, the data suggest application to determinations that require resolution of potassium and rubidium spectral emission lines. Results are immediately available from the strip chart, eliminating the usual time losses of exposure, development, and plate reading required for spectrographic procedures. The flame photometric technique for the determination of rubidium in thc presence of large
concentrations of potassium is easier and appears more accurate than flame spectrographic methods. LITERATURE CITED
(1) Glendening, B. L., Parrish, D. B., Schrenk, W. G., ANAL. CHEM. 27, 1554 (19%). \----,.
(2) Kingsley, G. R., Schaffert, R. R., Science 116, 359 (1952). (3) Pro, M. J., Nelson, R. A., blathers, A4.P.. J. Assoc. Ofic. - Aqr. Chemzsts 39,
506 (1956).
RECEIVED for review July 6, 1959. Accepted November 4, 1959. Contribution 586, Department of Chemistry, Kansas Agricultural Experiment Station, Kansas State University, Manhattan, Kan. Work supported in part by a grant from the U. S. Public Health Service.
Spectrophotometric Determination of Cobalt and Nickel with Oxamidoxime GEORGE A. PEARSE, Jr.,l and RONALD T. PFLAUM Department of Chemisfry, Sfafe Universify of Iowa, Iowa Cify, Iowa
b A method i s presented for the determination of cobalt and nickel in a mixture of diverse ions. Oxamidoxime is used as the reagent for the simultaneous precipitation of nickel and the complexation of cobalt as a soluble colored species. Nickel is quantitatively precipitated as Ni(C2HbN40212. 2Hz0 at pH 8 to 9.5. The formation of the yellow cobalt oxamidoxime complex is quantitative under identical conditions. The method for cobalt and nickel is rapid and accurate within spectrophotometric limits. Results on selected samples and the effects of diverse ions are discussed.
T
determination of cobalt and nickel has long presented a problem to the analytical chemist. As these two transition metals readily form colored coordination compounds with many organic ligands, the colorimetric method of analysis has been widely used. However, the reagents employed have not been specific for either metal ion, and the methods developed have suffered from interferences due to the presence of other metal ions (3, 4, 6, 9). A study of the reactivities of amidoximes with transition metal ions disclosed useful color reactions resulting HE
Present address, E. I. du Pont de Semoura & Co., Inc., Seaford, Del.
from the coordination of cobalt(I1) ion with oxamidoxime ( 7 ) . Precipitation of nickel(I1) ion with oxamidoxime, HON: (NH2)C-C(KHz):NOH, has been reported ( 1 ) and vias observed in this study. As a consequence, the cobalt(11) and nickel(I1) oxamidoxime systems were investigated thoroughly and a spectrophotometric method for the two metal ions was developed.
weighed amount of reagent in deionized water . Standard solutions of cobalt(I1) and nickel(I1) ions were prepared from the corresponding perchlorate salts (G. Frederick Smith Chemical Co.). Cobalt solutions were standardized gravimetrically using 3,5-dimethyIpyrazole as the cobalt precipitant (8); nickel solutions, by gravimetric determination of the dimcthylglyoxime precipitate. ,411 other reagents were prepared from reagent grade chemicals.
APPARATUS AND REAGENTS
All spectrophotometric measurements were made a t 25" C. with a Gary Model 11 recording spectrophotometer, using 1-em. matched silica cells. A Beckman Model G pH meter was used for all p H measurements. Oxamidoxime was prepared by the reaction of 0.5 mole of dithio-oxamide (rubeanic acid) with 1 mole of hydroxylamine (6). Hydroxglammonium chloride, neutralized with an equivalent amount of sodium carbonate, was added slowly, with stirring, to a hot solution of the dithio-oxamide in methanol. After refluxing for 0.5 hour on a steam bath, the solution was concentrated to one half its volume and cooled. The resulting crystals were recrystallized from water with the use of decolorizing charcoal. A 75% yield of pure white compound, with melting point of 2023" C. (literature value 202" C.), was obtained. A 0.1M stock solution of oxamidoxime, stable for more than six weeks, wa3 prepared by the dissolution of a
RECOMMENDED PROCEDURE
Dissolve the sample containing cobalt and nickel by appropriate means using a minimum of solvent. Evaporate t o 1-nil. volume and add 15 nil. of 12M hydrochloric acid. Place the solution on a 20 em. X 0.75 sq. em. column of Dowex 1-X8, 50 to 100 mesh, resin. Elute the column with 15 to 20 ml. of 4M hydrochloric acid. Adjust the acidity of the sample solution to pH 6 and the volume to about 50 ml. Add 10 ml. of 0.1M reagent solution for each 5 mg. of combined cobalt and nickel ion and sufficient solid sodium acetate to raise the pH to 8 to 9.5. Collect the nickel oxamidoxinie precipitate (which mag' contain insoluble hydrous oxides of diverse heavy metal ions) into a medium-porosity sinteredglass filter and wash with two 5-ml. portions of O.1M sodium acetate solution. Combine the filtrate and the wash solution, adjust to volume with 0.1M sodium acetate, and measure the absorbance a t 350 mu. Calculate the VOL. 32, NO. 2, FEBRUARY 1960
213
amount of cobalt from a previously prepared calibration curve. Dissolve the nickel precipitate in 10 ml. of 2M hydrochloric acid and dilute to 100 ml. with distilled water. Measure the absorbance a t 233 m* and calculate the amount of nickel from a previously determined calibration curve.
I
I 2 0 -
',
i
I
1
U
DISCUSSION AND RESULTS
Reactions of Cobalt(II) Ion with Oxamidoxime. T h e addition of a solution of oxamidoxime t o a solution of cobalt(I1) ion has no visible effect in the p H range from 1 to 7 . The formation of a yellow color increases in intensity with increasing pH. The colored species is stable in the presence of oxidants or reductants and is not extractable into water-immiscible alcohols or common nonpolar solvents. It is dependent upon cobalt(I1) ion, hydrogen ion, and reagent concentration. Spectrophotometric measurements on solutions containing a constant excess of reagent over cobalt ion (10 to 1) showed that maximum absorption a t 350 mp was obtained for solutions a t pH 8 to 9.5. The molar absorptivity of the complex was 8550. Measurements were made a t 350 mp, as the absorbance of the excess reagent is negligible a t this wave length. The absorption maximum for the cobalt complex occurs a t a shorter wave length and is masked by the absorption of the excess reagent. The absorption curves (Figure 1) are representative of the cobalt oxamidoxime complex. Mole ratio studies of the cobalt system showed that a tenfold excess of reagent was sufficient for maximum color formation; however, greater amounts do not interfere. The system obeys Beer's law in the concentration range of 1 to 50 p.p.m. of cobalt as determined on solutions containing varying concentrations of cobalt ion, a tenfold excess of reagent, and sufficient sodium acetate to give a p H of 8 to 9.5. Reactions of Nickel(I1) Ion with Oxamidoxime. The addition of a solution of oxamidoxime to a solution of nickel(I1) ion has little effect in the p H range of 1 to 6. An orange precipitate is formed upon the addition of sodium acetate, a n alkali metal hydroxide, or ammonia. Nickel is quantitatively precipitated as Ni(C2K~N4Oz)z.2Hz0 from a solution a t pH 8 to 9.5 in the presence of a 5% excess of the reagent. The nickel oxamidoxime precipitate is readily filterable onto a mediumporosity sintered-glass filter. Washing with two 10-ml. portions of 0.1M sodium acetate solution removes all traces of the original solution without loss of precipitate. The precipitate can be ovendried a t 110' C. and weighed as Ni(C2H6N*O2)e. A sample of the precipitate was analyzed for nickel by dwtructive oxidation of the organic ligand in
214
ANALYTICAL CHEMISTRY
250
350 WAVE
450 LENGTH
650
550
IMP)
Figure 1. Absorption spectra of cobaltoxarnidoxime system in sodium acetate solution, pH 9 1. 2.
I.
Table
Ion
2 X 10-4M cobalt 1 X 10-4M cobalt
Effect of Diverse Ions
c16;+S
c u +z l7-
Ag +
Na +
so,-2
Vfs
4.
4 X 10-6M oxomidoxime
I
i
I
I
I
Amt. Permissible, P.P.M. 2000 500" . ~
Ca +-2
3. 5 X 10-6M cobalt
.
2000
2000 2000 300" 0 1000 0 200" 2000 500" 1000 2000 500 2000 2000
1005 Zn +2 2000 Interference to Ni determination. Interference to Co determination.
acidic solution and precipitation of nickel as the dimethylglyoxime complex in the usual manner. Calculated for N ~ ( C Z H ~ N ~ ONi, ~ ) ~20.04%; : found, 19.86%. The water-insoluble nickel-oxamidoxime complex is readily soluble in dilute mineral acid, in strongly alkaline solution, and in dimethylformamide. An aqueous solution of the dissolved precipitate a t a pH of 3 to 5 shows an absorption maximum a t 233 m,u. This coincides with the wave length of maximum absorption for the reagent in acidic solution. Absorption curves for the nickel-oxamidoxime system are shown in Figure 2. The molar absorptivity for the protonated reagent is 10,500 a t the above wave length. The molar absorptivity, based upon nickel in the original precipitate, is 21,000. The nickel system obeys Beer's law in the concentration range of 1 to 30 p.p,m. of nickel. Effect of Diverse Ions. The effects
Figure 2. Ultraviolet spectra of nickel oxamidoxime system in acidic solution, pH 3.5 1. 2 X 10-'M oxamidoxime 2. 7.2 X 10-6M nickel 3. 3.6 X 10-6M nickel
of diverse ions on the cobalt and nickel oxamidoxime reactions are summarized in Table I. The data were obtained on solutions containing 5 p.p.m. each of cobalt and nickel ions and the particular diverse ion under investigation. Diverse cations were added as the chloride or perchlorate salts. Alkali metal salts of the anions were used. Iron and copper react directly with the organic reagent and cause serious interference. These ions, together with zinc, can be easily removed by the use of an anionic exchange resin. The technique developed by Kraus (5) for the separation of the first series of transition metal ions via the interaction of their chloro complexes with an anionic resin proved to be particularly applicable to the present problem. Insoluble hydrous oxides formed with heavy metal ions can be removed, together with the nickel complex, from the soluble cobalt complex by filtration. Large amounts of these gelatinous precipitates interfere in the quantitative recovery of the nickel complex. However, after dissolution in
acidic medium, the prcwnce of heavy metal ions does not interfere in the subsequent measurements. Silver ion consumes reagent through a redox reaction and vanadium reacts to give a colored coordination compound. Results on Selected Samples. T h e results of determinations of cobalt and nickel on selected samples are presented in Table 11. The synthetic samples were prepared from mixtures of t h e perchlorate salts. Kational Bureau of Standards sample 161 n.as selected for its relatively low percentage of iron. Cobalt and nickel can be determined in varying ratios with the usual spectrophotometric accuracy of 1 2 7 , . Cominon anions and other metal ions, n-ith th(. pxception of those described above, do not offer serious interferencc in the concentration range of 500 to 2000 p.p .m. ACKNOWLEDGMENT
The authors thank the Research Corp.
Table 11.
Summary of Determinations
Cobalt, Present 6.25 SI 2.60 S? 1.25 Sa S4(20 nig. Fe 20 mg. Zn) 2.75 Sj(10 mg. AI 10 mg. Cu 20 mg. Fe) 5.80 S6 (10 mg. Cu 20 mg. Fe 20 mg. Zn) 2.60 NBS 161 Ni-Cr casting alloy 0.23 Sample
++
+
++
for the financial assistance which made this work possible. The authors also thank the hlallinckrodt Chemical Works for the generous sample of dithiooxamide. LITERATURE CITED
(1) Chatterjee, R., J. Indian Chem. Soc. 15. 608 (1938). (2) Fische;, E.,'Ber. 2 2 , 1931 (1889). (3) Gupta, H. K. L., Sogani, N. C., ANAL. CHEM.31, 918 (1959). (4) Jonassen, H. B., Chamblin, V. C., Wagner, V. L., Jr.,lbid., 30,1660 (1958).
Mg. Found 6.19 2.60 1.24 2.72 5.86
2.53 0.22
Kickel, Mg. Present
Found
1.25 2.60 6.25 4.55 5.80
1.26 2.54 6.11 4.50 5.80
2.60
31.50
"57
31 , 1 5
(5) Kraus, K. -4.,Moore, G. E., J . -4m. Chem. SOC.75, 1460 (1953). (6) McDowell, B. L., Meyer, A. S.,Jr., Feathers, R. E.,
ANAL.CHEM.31, 931 ( (7) Pearse, G. A., Jr., Pflac J . Am. Chem. Soc., in press. ( 8 ) Pflaum, R. T.. Dieter. L. H.. Proc. Iowa Acdd. Sci. 64, 235 (1957). ' (9) Wheatley, R. D., Colgate, S. O., ANAL.CHEM.28, 1897 (1956). '
RECEIVED for review March 17, 1959. Accepted November 5, 1959. Division of Analytical Chemistry, 135th Meeting, .4CS, Boston, Mass., April 1959.
Determination of Iron and Iron-Aluminum Mixtures by Titration with EDTA DONALD G. DAVIS' and WILLIAM R. JACOBSEN2 School o f Chemistry, Georgia institute o f Technology, Atlanta 7 3, Ga.
b Iron in the presence of aluminurn has been titrated with EDTA. To avoid the interference of aluminum, a pH of 1.0 was used. The end poin', was determined spectrophotometrically with 5-sulfosalicylic acid acting a the indicator. Results were accurate to 0.3%, even though the amount of aluminum was twice that of iron, if the iron concentration was not much less than 10-'M in the solution being titrated.
S
methods for the direct titration of iron with (ethylenedinitrilo) tetraacetic acid (EDTA) have been proposed ( I O ) . Although Cheng, Bray, and Kurtz (3) have reported that aluminum does not interfere with the titration of iron at p H 2 t o 3 except in "large amounts," Sweetser and Bricker ('7) found t h a t the presence of aluminum caused noticeable EVERAL
1 Present address, Louisiana State University in New Orleans, New Orleans, La. Present address, Savannah River Plant, E. I. du Pont de Nemours & Go., Inc., Aiken, S. C.
positive errors in the p H range between 1.7 and 2.3. Because the findings of Sbveetser and Bricker were confirmed in this laboratory, a theoretical and experimental study was undertaken to establish conditions under which the titration of iron could be performed free from the interference of aluminum. CHEMICALS AND APPARATUS
,411 chemicals used were the best obtainable, usually reagent grade. Standard solutions of iron and aluminum rvere prepared by dissolving a weighed amount of the pure metal in the appropriate acid and diluting t o exactly 1 liter. The standardization of the iron solution was checked by titration with potassium dichromate solution according t o the classical procedure. The titrant solutions of EDTA were prepared by dissolving the reagent grade disodium salt of (ethylenedinitri1o)tetraacetic acid in distilled water. These solutions were stored in a polyethylene bottle and standardized against a standard bismuth solution at p H 1.2 t o 1.4 (9). Bismuth was chosen as a standard so that the EDTA could be standardized at approximately the pH at which it was t o be used. This procedure eliminates slight errors due to
the presence of extraneous metal ions in the distilled water, such as calcium and copper. Ascorbic acid was added t o mask iron impurities in the bismuth. The indicator solution was a 2 weight yo solution of Eastman White Label 5-sulfosalicylic acid dissolved in mater. It was stable indefinitely. The spectrophotometric titrations were performed using a Beckman Model D U spectrophotometer. The cell compartment of this instrument was replaced by a 3.75 X 3.75 X 5.50 inch black painted aluminum box, with holes in the side to alloiy the passage of light from the optical system to the phototube. The buret tip protruded through the lid of the box into the titrate, which was contained in a 100-ml. tall-form beaker held in the light path. The solution was stirred between additions of titrant by a magnetic stirrer placed beneath the titration compartment. Calibration volumetric glassware \vas used throughout this work. p H was adjusted by using a Beckman Zeromatic p H meter equipped with the usual glass and calomel electrodes. PROCEDURE
An accurately measured portion of VOL. 32, NO. 2, FEBRUARY 1960
215
