drogen ion and not the chloride ion.
A plausible explanation is t h a t the hydrogen ion competes more favorably than the quaternary ammonium ion for the anionic metal complex a t higher hydrochloric acid concentrations. These hydrogen-associated complex anions have their effective negative charge lowered by one or more depending upon the number of hydrogen ions 15 hich associate. These weak complex acid species can associate with fewer or no quaternary ammonium ions, and the organic solubility of the metal complex decreases; hence, the distribution ratio of the metal decreases. Further Li-ork is in progress on the quantitative explanation of this phenomenon. The distribution curves of hafnium (IV), tantalum(V) , and molybdenum (VI), although characteristic for the metal ions, will not be discussed because these metals form fairly complicated oxyanionic complexes even in concentrated hydrochloric acid. If one takes the concentration of the ligand, which gives 50% extraction, as a criterion of the stability of the various cobalt( 11) complexes, i t is apparent from Figure 1 t h a t the stability of the complexes is in the order SCN- > NOz- > Br- > C1-, which is in agreement with reported stability constants (1). The reason why the nitrocobaltate(11) complexes do not extract as well as the bromo- and chlorocobaltate(I1) complexes must be due to the high oxygen content of the complex. Thus, the nitrocobaltate(I1) complexes are more “water-like” and less soluble in the organic phase. A similar argument may explain why the very stable (ethy1ene)dinitrilotetraacetate complex of cobalt(I1) extracts as poorly as i t does.
A comparison of the characteristic curves of all the metals studied in this paper in Figure 2 with those studied by Kraus et al. (9) suggests t h a t the mechanism of distribution by solvent extraction with quaternary ammonium halides is similar to the mechanism of distribution by elution chromatography with anion exchange resins. Thus, quaternary ammonium halides can be called “liquid anion exchangers.” Also, qualitative predictions about separations in one system should apply to the other system. I n general, the values of D, are lower in solvent extraction than in anion exchange chromatography. This will require multiple contactings in solvent extraction and frequent back extractions to obtain quantitative separations; however, this should not prohibit solvent extraction by quaternary ammonium halides from becoming an analytically useful technique where large capacity is of paramount importance. ACKNOWLEDGMENT
A portion of this work was supported through Grant No. At-(11-1)-775 of the U. S. Atomic Energy Commission. Paul Hovsepian gratefully acknowledges fellowship aid under the Summer Fellowship Program of the National Science Foundation for the summer of 1960. LITERATURE CITED
(1) Bjerrum, J., Sillen, L. G., Schwartzenbach, G., “Stability Constants. Part 11. Inorganic Ligands,” pp. 39, 97, 112, The Chemical Society, London, 1958. (2) Blaedel, W. T., Knight, H. T., ANAL. CHEM.26, 741 (1954). (3) Clifford, W. E., Bullwinkel, E. P., hlcClaine, L. A., Noble, P., Jr., J. Am. Chem. SOC.80, 2959 (1958).
(4) Coleman, C. F., Brown, K. B., Moore, J. G., Crouse, D. J., Ind. Eng.Chem. 50, 1756 (1958). (5) Fieser, L. F., “Experiments in Organic Chemistry,” p. 288, Heath and Co., Boston, 1957. (6) Good, M. L., Bryan, S. E., Abstract, pp. 43N44?i, 138th Meeting, ACS, Kew York, September 1960. (7) Good, M . L., Bryan, S. E., J . Am. Chem. SOC.82, 5636 (1960). (8) Katekaru, J., Freiser, H., abstract, p. 14B, 138th Meeting, ACS, New York, September 1960. (9) Kraus, K. A., Nelson, F., “Anion Exchange Studies of the Fission Products,” Vol. 111, Proceedings of the International Conference on Peaceful Uses of Atomic Energy, Geneva, 1955, p. 113, U. S., New York, 1956. (10) Lindenbaum, S., “Liquid Ion Eschangers,” Gordon Research Conferences on Ion Exchange, Tilton, ?;. H., June 1961. (11) Maeck, K. J., Booman, G. L.9 Elliot, M. C., Rein, J. E., ANAL. CHERI.30, 1902 (1958). (12) Metcalfe, L. D., Ibzd., 32, 70 (1960). (13) Moore. F. L.. Bull. KAS-XS-3101. ‘ 6ffice of’ Technical Services, Depart: ment of Commerce, Washington 25, D. C., 1961. (14) Morrison, G. H., Freiser, H., “Solvent Extraction in Analytical Chemistry,” Wiley, Sew York, 195:. (15) $fusser, D. F., Krause, D. P., Smellie, R. H., Jr., U. S. Atomic Energy Comm. Rep., AECD-3907. (16) Rosen, RI. J., Goldsmith, H. -4., “Systematic Analysis of Surface-Active Agents,” p. 10, Interscience, Sew York, 1960. (17) Schindewolf, L.> Progrcss Report, p. 23, RIass. Inst. Tech. Lab. for Xuclear Science, Feb. 23, 1956. (18) Silverman, L., LIoudy, I., Hawley, D. W., A s 4 ~ CHEV. . 2 5 , 1369 (1953). (19) Smith, E. L., Page, J. E., J . SOC. Cheni. Ind. (London) 67, 18 (1948). RECEIVED for review June 5 , 1961. -4ccepted Sovember 13, 1961. 8th Anachem Conference, Detroit, Nich., October 1960.
Use of Manganese(ll1) for Spectrophotometric Determination of Ferrocyanide, Tin(ll), and Iron(ll) WAHID U. MALIK and MOHAMMAD AJMAL Department o f Chemistry, Aligarh Muslim University, Aligarh, India
b Investigations of the oxidimetry of Mn(lll) have revealed that manganic sulfate can be employed for determination of the ferrocyanide (0.000804to 0.00296M), tin(ll) (0.001 to 0.01 3M), and iron(ll) (0.001 to 0.0999M) by carrying out absorption measurements at 5 2 5 my (absorption maximum for manganic sulfate solution). Better results for the estimation of ferrocyanide are, however, obtained b y working at 425 mp (absorption maximum for potassium ferricyanide). Nitrite, SUIfite, and iodide interfere with the determinations.
M
aspects of the chemistry of Mn(II1)-viz., the use of the reagent as an analytical oxidant, its complexes, the influence of foreign ions on its oxidation potential, etc.-have not been fully investigated and need a more systematic study. Ubbelohde (4) pointed toward its possible use a s a n analytical reagent in volumetric analysis. Recently Malik (2) studied the reaction between Mn(II1) and potassium ferrocyanide and ascertained the following: Mn(II1) oxidizes potassium ferrocyanide to potassium ferricyanide, the change being possible in view of ANY
the high oxidation potential ( - 1.51 volts) of the couple hIn+3 e Mnf2; in a dilute solution of potassium ferrocyanide a perceptible change in color (from colorless to yellow) is observed on the addition of manganic sulfate solution; mith concentrated solutions of the reactants the oxidation of potassium ferrocyanide t o the ferricyanide is followed by the mutual interaction between hln(II1) and the oxidized product, yielding a n insoluble complex of the composition : hfn( 1II)Fe (111) Cyg. These observations led us to in-
+
VOL. 34, NO. 2, FEBRUARY 1962
* 207
vestigate the use of lIn(l11) in osidimetry for the estimation of ferrocyanide, tin (11), and iron(I1). The work described here includes the possible interference of certain anions like nitrite, sulfite, and iodide in these estimations. EXPERIMENTAL
Reagents. hlanganic sulfate solution \$as prepared by the method recommended by Ubbelohde ( 4 ) , except t h a t a larger amount of sulfuric acid was added t o keep the solution stable. I t s strength was determined by adding a knou-n amount of ferrous aninionium sulfate a n d then titrating the ewess ferrous ammonium sulfate against potassium permanganate of knon n strength. An approximately 0.551 potassium ferrocyanide solution was prepared by dissolving the recrystallized product in doubly distilled water. The solution was stored in dark bottles and the strength determined by titrating against potassium permanganate ( 3 ) . Six grams of stannous chloride (British Drug Houses) was dissolved in
Table 1. Determination of Potassium Ferrocyanide, Stannous Chloride, and Ferrous Ammonium Sulfate with Manganic Sulfate
Manganic Sulfate
Aliquot from Concn. Taken, Curve, Found, 'If M1. Ml. I%! Potassium Ferrocyanide with 0.01M Manganic Sulfate 0.00296 1.0 0.30 0,003 1.5 0.45 0.00294 2.0 0.59 0,00295 0.0060 3.0 0.60 0.0060 4.0 0.80 0.006 5.0 1.0 0.06 0.000804 3.0 0.24 0.0008 5.0 0.40 0.0008 7.0 0.56 0,0008 Stannous Chloride with 0.065M Manganic Sulfate 0.013 5 1.0 0.013 7 1.5 0.013 8 1.7 0.013 0.0307 5.0 2.35 0.305 8.0 3.80 0.308 10.0 4.80 0.304 Kith 0.01M Manganic Sulfate 0,001 5 1.0 0.001 10 2.0 0.001 Ferrous Ammonium Sulfate with 0.076M Manganic Sulfate O.OH99 1.0 1.3 0.1 2.fi 0.1 2.0 3.9 0.i 3.0 2.0 0.05 1.3 0.05 2.6 4.0 0.05 3.9 6.0 0.05 With 0.01M Manganic Sulfate 0,001 5.0 1.0 0.001 10.0 2.0 0.001 Concn. Taken,
208
ANALYTICAL CHEMISTRY
0 2
04
0 6
08
IO
'dot M a n s a . ~ c s u m r
Figure 1 . A.
12
14
, nr
Titration curves RESULTS A N D DISCUSSION
Ferrocyanide with manganic sulfate at 425
mp
6. Ferrocyanide with manganic sulfate a t 5 2 5 m!J
C.
A rather abnormal behavior is seen in experiments with sodium nitrite. The presence of even small amounts of this salt (1 cc. of 0.0lM) instantaneously increases the absorbance, the value being the same (1.5) a t 425 mp as when equivalent manganic sulfate had been added t o potassium ferrocyanide solution. A few results are given in Tables I and I1 and typical curves are shocvn in Figure 1.
Tin(ll) or iron!ll) with manganic sulfate at
525
mp
100 ml. of concentrated hydrochloric acid and the total volume mas made up to 1 liter. The strength of the solution was determined iodometrically ( 5 ) . The solutions of sodium nitrite (B.D.H.), sodium sulfite (B.D.H.), and potassium iodide (analytical reagent grade, B.D.H.) mere obtained by dissolving known amounts of the reagents in doubly distilled water. The solution of manganic sulfate was diluted with 1 0 s sulfuric acid to check hydrolysis, and other solutions were diluted with air-free distilled r$ater. Procedure. The technique used for the determinations was that employed by Bricker and Loeffler ( 1 ) in cobaltic osidimetry. The absorbance measurements were carried out with a Bausch & Lomb Spectronic 20 colorimeter employing the 1/2-inch tube. A fixed amount of the reagents to be determined was taken in the tube and varying amounts of iLIn(II1) r e r e added. The effect of foreign ions wis studied by determining the absorbance of a mixture containing varying amounts (0.2, 0.4, 0.6, 0.8, 0.9, and 1.1 ml.) of 0.01M manganic sulfate, fixed amounts (2 ml.) of potassium ferrocyanide (O.O3M), iron(I1) ( O . O O l X ) , and tin(I1) (O.OOlM), and 1 ml. of the interfering ions (O.OlM), making the total volume 5 ml. I n the case of iron(I1) and tin(I1) the break in the absorbance curves is realized a t a much later stage (in the presence of nitrite and sulfite) because of the simultaneous oxidation of the reagents and the added ions. T i t h potassium iodide the behavior was a little different. Here the absorbance values increase continuously, since the liberated iodine also gives absorbance a t 525 mp. Sulfite ions do not interfere with ferrocyanide determination when present in very small quantities (0.001M). However, in the presence of potass'mm iodide and higher concentrations of sodium sulfite interference does occur, the absorbance value increasing markedly in the case of potassium iodide.
The values of K (calculated from the oxidation potentials) for the reactions F e c ~ , -s~ Mn+* F e c ~ , - ~ , 2AIn+3 Snf2 + 2Mn+2 Sn+4, and Fe+2 z Fe+3 hIn+2 are 3.09 X lo", 11.2 X lo2?,and 20.99 X lo", respectively, showing that ferrocyanide ions are completely oxidized to ferricyanide ions in the presence of hfn(II1). Colorimetric estimations can, therefore, be carried out a t the absorbance maximum of RIn+3and also at the absorption maximum of F e c ~ , - ~ (in the case of potassium ferrocyanide estimations). Both for tin(I1) and iron(I1) measurements were carried out at 525 mp [maximurn absorption for Mn(III)] and sharp breaks were realized after the complete osidation of these two ions (Figure 1, C). I n the potassium ferrocyanide determination the end point was sharper when the titrations were carried out a t 425 mp (maximum absorption for FeCy6+, than a t 525 mp (Figure 1, A , B ) . Bricker and Loeffler (1) estimated Fe-2 a t 350 m p . We could not carry out these measurements because of high absorbance of AIn(II1) solution a t this m-avelength. The experimental results given above provide evidence that SIn(II1) can be successfully employed for the determination of ferrocyanide ions (0.000804 t o 0.00296M), stannous ions (0.001 to O.O13M), and ferrous ions (0.001 to O.O999M), provided such interfering ions as iodide, nitrite, and sulfite are not present. I n most cases
+ + +
Table It.
+ + +
Interference in Ferrocyanide Determination
Vol. of 0.03M K4Fe(CN)G 2 ml. Vol. of (0.01M) interfering salt 1 ml. 0.01M ManAbsorbance ganic Sulfate, Dil. hI1. Na2S0p KanSOaNitrite Iodide 0.51 0.34 1 . 5 0.68 0.2 0.4 0.95 0.61 1 . 5 1 . 2 0.80 1 . 5 1.6 0.6 1.3 1.5 1.8 1.5 1.2 0.8 1.3 1.5 2.0 0.9 1.5 1 . 5 Beyond 1.5 1.4 1.1 scale a Dilute Na2SOa, 0.001M.
even traces of these foreign ions interfere with the results.
one of them (M.A.) to carry out this work.
ACKNOWLEDGMENT
LITERATURE CITED
The authors are grateful to A. R. Kidwai, head of the department, for providing facilities and to C.S.I.R. (India) for awarding a fellowship to
(1) Bricker, ’’9
c. E., Loeffler, L. J.,
1418-23 (1955)*
( 2 ) Malik, W. U., J. Indian Chem. SOC. 38,303-12 (1961). (3) Sutton, F., “Volumetric A4nalysis,”
10th ed., p. 217, Churchill, London,
1911.
(4) Ubbelohde, A . R. J. P., J. Chem. SOC. 1935, 1605-7.
(5) Vogel, A. I., “Textbook of Quantitative Inorganic Analysis,” 2nd ed., p. 352, Longmans, Green, New York, 1959.
RECEIVEDfor review May Accepted November 16, 1961.
8, 1961.
Spectrophotometric Determination of Boron Using Barium Chloranilate RAM D. SRIVASTAVA, PAULA R. VAN BUREN, and HYMAN GESSER Parker Chemistry Laboratory, University o f Manitoba, Winnipeg, Manitoba, Canada
b Boron as boric acid has been determined spectrophotometrically b y reacting it with tartrate reagent and barium chloranilate. The method i s based on the subsequent liberation and analysis of colored acid chloranilate ion. The procedure has shown accuracy and precision to h O . 1 pg. of boron over the range of 0 to 10 pg. of boron. The method is simple, sensitive, and better than some of the photometric methods used in boron analysis.
T
HE USE of barium, mercury(II), strontium, lanthanum, and thorium chloranilntes as analytical reagents for the photometric determination of anions has been reported only within the past few years. The method is based on the metathesis of chloranilic acid salts to form an insoluble salt of the anion to be estimated and the subsequent liberation of highly colored acid chlornnilate ion. Such a method has been successfully applied for the estiniation of sulfate ( d i , chloride ( 1 ) . and fluoride ( 7 ) . I t is realized that the method is not limited to the determination of sulfate, chloride, and fluoride only, but can be extended to various other anions if a suitable salt is found which gives an insoluble precipitate with the anion to be investigated. Recently, Hayashi, Ihnzuka, and Ueno ( 6 ) have determined 3 to 300 p.p.m. of phosphate using lanthanum chloranilate. Various salts of chloranilic acid such as barium, mercuric, lanthanum, and thorium were tried for the estimation of boron, without success. Nercuric chloranilate gave a precipitate with sodium borate but not with boric acid. This precipitate was considered to be that of mercuric oside rather than of mercury borate. Apparently, borates of these metals are not formed under the conditions of observation, because no acid
cliloranilate ion is liberated when boric acid is added to the metal chloranilate. Uoron as boric acid is quantitatively precipitated as a comples borotartrate (4) with tartaric acid reagent and barium chloride. Attempts were, therefore, made to estimate boron as boric acid by reacting it with tartrate reagent and barium chloranilate to form a complex borotartrate and the acid chloranilate ion. The amount of the latter formed should be proportional to the original amount of boron present. PROCEDURE
To a n aliquot of boric acid containing up to 100 pg. of boron in less than 4 ml. in a IO-ml. (boron-free) volumetric flask are added 1 ml. of tartaric acid reagent buffer (1.4 grams of tartaric acid, 24 grams of ammonium chloride, and 10 ml. of 15W ammonium hydroxide in a totnl volume of 100 ml.) and 5 ml. of acetone. The mixture is diluted to volume with distilled water, approximately 0.04 grani of barium chloranilate (Fisher Scientific Co.) is added, and the flask is shaken intermittently for 30 minutes. The excess barium chloranilate and the precipitated barium borotartrate are removed by filtering through a Whatman S o . 44 filter paper. The pH of the final solution was 8. The absorbance of standard boron solutions is measured against a reagent blank a t 530 mp in a 5-em. cell using a Coleman 14, Universal spectrophotometer. The results are given in Table I. RESULTS A N D DISCUSSION
The sensitivity of the method can be increased by using the near-ultraviolet absorption band of the acid chloranilate ion. In this region, the absorption of chloranilate solutions shows a sharp peak which is about 20 to 30 times more intense than the absorbance in the region of 530 to 560 mp. The position of the peak varies (5) with the p H and the nature of the solvent system.
According to Bertolacini and others (S), the barium and strontium chlor-
anilates react with the p H 9 buffer, to produce some substance that absorbs strongly a t 300 to 335 mp. They, therefore, preferred to determine sulfate and fluoride a t a p H of 4, thereby sacrificing the sensitivity slightly. The present method for the determination of boron works only in alkaline medium. The ammonium hydroxide-ammonium chloride buffer of p H 8 ivhich was used also reacted with barium chloranilate. The observations were, therefore, made a t 355 mp which is outside the range of this interference.
Table
1.
Absorbance of Boron Solutions
Standard
(Coleman Model 14 Universal spectrophotometer, 530 mp, 5-cm. path length) P.P.M. of Boron Absorbance 0 3 0 01 0 5 1 0
1.5 2.0 3 .O 4.0
5.0
7.0 8.0
10.0 Table
II.
0 09 0 13 0 16 0.19
0.26 0.28 0.35
0.41 0.42 0.47
Absorbance of Standard Boron Solutions
(Beckman DB spectrophotometer, 355 mp, 1-cm. pat.h length) P.P.M. of
Boron
Absorbance
VOL. 34, NO. 2, FEBRUARY 1962
209
