V O L U M E 26, NO. 2, F E B R U A R Y 1 9 5 4
375 Evans, D. P.. and Simmons, N. T., J . Soc. Chem. I n d . (London),
stock standard solutions with distilled water through which carbon dioxide had been passed for 20 minutes. Carbon dioxide was also passed through the other solutions used in these titrations and the titrations were done under carbon dioxide. These titration values are slightly high, and show an absolute error of 3 y with quantities of iron greater than 500 y and of less than I y with smaller quantities. The data from the dichromate titrations indicate that under the conditions of these titrations the errors are within those of the experimental measurements.
63. 29 f19441,
Harris, E: D., and Lindsey, d.J., Analust, 76, 647 (1951). Hewson, G. W., and Rees, R. L., J . Soc. Chem. Ind. (London), 54, 2 5 4 T (1935).
Kolthoff, I. XI., Z . anal. Chem., 6 0 , 341 (1921). Meier, D. J., Myers, R. *J., and Swift, E. H., J . A m . Cheni. Soc., 71, 2340 (1949). Ramsey, W. J., Farrington, P. S., and Swift, E. H., AXAL. CHEM.,22, 332 (1950).
Reith. J. F.. Biochena. Z.. 216. 249 11929). Sadusk, J. F., and Bell, E. G., IND.ENG CHEW.,SAL. ED.,5 ,
ACKNOWLEDGMENT
3 5 6 (1933).
Topf, G.. Z . anal. Chem., 26, 137 (1887). Tutrundic, P. S., and Mladenovic. S.,A d . Chzm. Acta, 8 , 184
The authors are indebted to Jerry C. Mitchell for aid in carrying out the confirmatory titrations of thiosulfate.
(1953).
LITER4TURE CITED
R E C E I V Sfor D reriew August 2-1, 1953. Accepted October 2 3 , 1953. PreCHEMICAL SOCIETY,Los s m t d a t the 123rd Meeting of the AMERICAN
(1) Bassett, H., and Durant, R. G., J . Chena. S o c , 1927, 1458. (2) Bradbury, J. H., and Hambly, A. K.,Australian J . Sci. Research. A5, 541 (1952).
Angelen. Calif., hlarch 1953. Contribution KO.1841, Gates and Crellin Laboratories of Cheniistry, California Institute of Technology, Pasadena, Calif.
Acetylacetone as an Analytical Extraction Agent Extraction of Aluminum, Gallium, and Indium JOHN
F. S T E I N B A C H and H E N R Y FREISER
D e p a r t m e n t o f Chemistry, University o f Pittsburgh, Pittsburgh 73, P a .
-is acetylacetone has already proved to be a useful and interesting extraction agent, the extraction characteristics of some trivalent metals, such as aluminum, gallium, and indium were studied. Equilibrium extraction curves for these metals by acetylacetone have been determined. The aqueous solubility of acetylacetone as a function of pH and ionic strength has been determined. The relation between the shape and position of the extraction
K
OLTHOFF and Sandell ( 6 ) and Irving and Williams ( 6 ) have derived the following equation for the distribution of metal chelates between organic and aqueous phases:
n here p , and p , are the pal tition coefficients of the chelate and ieagent hehveen the tn-o phases; K f is the over-all formation constant of the chelate, JIR,, and K , is the acid dissociation constant of the reagent, HIt. The subscripts o and zu refer to the organic and water phases, respectively. Usuallj- the solubility of the chelate is so much greater in the organic phase than it is in water that the term ( p c - D )reduces to p,. From this equation it can be seen that the distribution is dependent on two variables, ( H +) and (HR),. Consequently, two methods of verification may be employed. The p H may be maintained constant through use of a buffer and the distribution Ftudied as a function of reagent concentration. However, as pointed out by Irving ( 5 ) , more useful analytical data may be obtained by maintaining the reagent concentration constant and studying the distribution as a function of pH. Irving has shown that when the per cent extracted is plotted against the p H when the reagent concentration is constant, a sigmoid extraction curve whose shape is independent of its position along the p H axis is obtained. The slope of the curve depends only on the valence of the metal; the greater the valence, the greater the slope. The p H of 50% extraction, ( D = I ) , has been called the PHI/*value and the spread of these values for two metals is an indication of their separability in extraction processes.
curves and properties of the acetylacetone chelates is examined and the usefulness of such curves in predicting the feasibility of analytical separations and determinations is indicated. The factors which must be controlled in obtaining equilibrium extraction curves are discussed. The results indicate the possible separations of aluminum, gallium, and indium and have a more general bearing on the development of solvent extraction as an analytical tool.
However, it is rather difficult t o maintain a constant reagent concentration in the organic phase when a reagent such as dithizone is used in, for example, a chloroform-water system, because the distribution of the reagent itself is a function of p H and also the amount of reagent consumed in the formation of the chelate must be considered if the ratio of metal to reagent is large. If acetylacetone is used as both the solvent and reagent, the concentration of acetylacetone in both organir and aqueous phase is of necessity constant, providing the solubility of acetylacetone in water does not vary appreciably n i t h pH. Hence, equilibrium extraction curves are more easily obtained. Other advantages of acetylacetone as a solvent have been discussed ( I O ) . In this paper the extraction behavior of aluminum, gallium, and indium is considered, along with the solubility of acetylacetone in water as a function of pH which must be knoxn to evaluate the term, (HR),. EXPERIMENTAL
Purification of Acetylacetone. The purification of the acetylacetone has been discussed (IO). Solubility of Acetylacetone in Water as a Function of pH. Since existing methods ( 8 , 9 ) of determining acetylacetone in Rater were found unsatisfactory when excess acid was present, a method of analysis was developed in which the amount of chelate formed on adding an excess of metal ion to the buffered sample would indicate the amount of acetylacetone originally present. An excess of ferric ion was added to a I-ml. sample of dilute acid saturated with acetylacetone. Sufficient strong sulfuric acid-sodium sulfate buffer was added to bring the p H of all
ANALYTICAL CHEMISTRY
376
samples to 0.55. The extinction of the resulting mixture as measured a t 700 mfi (slit width = 0.04 mm.) on a Beckman Model DU spectrophotometer was employed as an indication of the amount of acetylacetone present. Beer’s law was obeyed up to a concentration of 0.300 gram of acetylacetone per ml. of solution. Analytical results (Tables I and 11) showed that acetylacetone was soluble to the extent of 17.0 grams per 100 ml. of solution a t a pH of 6.5 and remained a t this value down to a pH of 1. Below pH 1 there was a gradual decrease in solubility, probably caused by the salting-out effect of the larger concentrations of sulfuric acid. Preparation of Acetylacetonates. ALUMINUIIACETYLACETONATE. Reagent grade aluminum.nitrate contained a trace of iron which caused the aluminum acetylacetonate prepared when acetylacetone was added to an ammoniacal solution of aluminum nitrate to be orange-colored, presumably because of coprecipitation of a small amount of ferric acetylacetonate. Since the colored impurity interfered with the analysis of aluminum in extraction processes, efforts were made to prepare a colorless aluminum acetylacetonate. Repeated recrystallization of the product did not completely remove the color. When an aluminum nitrate solution whose pH was adjusted to about 0.5 with sulfuric acid was extracted with acetylacetone, the yellow compound was extracted into the acetylacetone layer while the aluminum remained essentially in the water phase. A solution of 150 ml. of mater and 10 grams of aluminum nitrate was extracted with three successive portions of acetylacetone. When the extracted water phase was treated with I S ammonium hydroxide, pure white aluminum acetylacetonate u-as preciDitated. On recrystallization from 95% alcohol colorle~s piisms were obtained. Analysis for AI(C5H,02)3, theoretical: 55.55% carbon, 6.53 hydrogen; found: 55.63% carbon, 6.44% hydrogen. GALLIUM Gallium metal was washed with . ACETYLACETONATE. warm 6.V nitric acid to remove traces of iron. It was then dissolved in concentrated nitric acid containing 10% by volume of sulfuric acid. When acetylacetone and base were added a white Figure 1. Distribution of Aluminum, Gallium, and precipitate of gallium acetylacetonate was formed. I t was recrystallized from 95% alcohol. Analysis for Ga(C5H702)3, Indium between Acetylacetone-Water Solutions of Varying pH theoretical: 49.21% carbon, 5.76% hydrogen; found: 49.38% carbon, 5.80% hydrogen. IXDIUM ACETYLACETONATE. The usual method ( 7 ) of preparing indium acetylacetonate gave a red-brown product but, beExtraction Curves. The distribution of aluminuni, gallium, cause of the chemical similarity of aluminum, gallium, and inand indium between acetylacetone and water solutions of varying dium, it was assumed that indium acetylacetbnate should be also pH I\ as determined using the general extraction procedure outcolorless. The preparation of indium acetylacetonate by the lined previously (10). The method of analysis which was emacetylacetone extraction procedure gave colorless crystals of indium acetylacetonate. A microanalysis showed 43.30y0 cartmn ployed is given below. and 5.13% hydrogen, while the theoretical analysis for indium The same curve for aluminum ivas obtained with the aluminurn acetylacetonate is 43.69% carbon and 5.14% hydrogen. initially in the acetylacetone layer or initially in the aqueous phase. Also the same curve was obtained using sulfuric, nitric, 01 hydrochloric acid to acidify the aqueous phase. Above a p H of 3.5, however, it was difficultto obtain equilibrium values of the Table I. Solubility of Acetylacetone as a Function of pH per cent extracted. The results of the extraction studies are and Temperature qhown in Figure 1. Solubility of Absorbanre Acetylacetone, .ANALYSIS OF ALUMIKUM,GALLIUU, AND INDIUM. Five-niillia t 700 mp G /MI. liter samples of the acetylacetone layer were removed and placed SCLFCRICACIDUSEDTO ADJUSTp H , 2 5 O C in 6-inch test tubes. Since acetylacetone will extract traces of 0.440 6.22 0.173 acid from the water layer, a few milligrams of sodium bicarbonate 0.440 4.12 0.173 0.439 2.61 0.173 were added to each test tube. Otherwise the acid traces would 0.435 2.00 0.170 ‘ cause an interfering coloration of the oxine later added as a colori0.430 1.49 0.168 0.430 0.99 0.168 metric reagent for the metals. The test tubes containing the 0.420 0.60 0.160 0.390 0.03 0.143 samples and sodium bicarbonate were allowed to stand overnight, Then 2 ml. of a 1% solution of oxine in alcohol were added and the SULFURIC ACIDUBEDTO ADJUSTpH, 40° C. mixture was thoroughly stirred. Oxinates are more stable than 0.470 4.62 0.193 0 469 3.25 0.193 acetylacetonates and oxine will convert them to oxinates which 0.469 3.25 0.193 have an absorption peak a t 395 mp. Beer’s law is followed by 0.470 2.65 0.193 0.472 2.08 0.194 solutions of aluminum in acetylacetone containing up to 0.4 gram 0.470 1.62 0.193 1.oo 0.465 0.190 of aluminum per liter. Solutions containing 0.001 gram of alumi0.469 0.50 0.186 num per 10 ml. were used for the aluminum extraction curves 0.443 0.13 0.175 determined here. PERCHLORIC ACIDUSEDT O -%DJWST pH, 25’ C. Both indium and gallium give more sensitive reactions with 5.55 0.440 0.173 oxine, and solutions containing 0.0003 gram of their acetylace4.91 0.439 0.173 3.00 0.442 0.173 tonates per 10 ml. of acetylacetone were employed. 2.53 0.440 0.173 2.21 1.87 1.55 1 .oo
0.46
0.445 0.444 0.449 0.450 0.485
0.176 0.175 0,180 0.186 0.200
DISCUSSION
The gradual decrease in the solubility of acetylacetone below pH 1 may be attributed to the salting-out effect of the higher
V O L U M E 26, NO. 2, F E B R U A R Y 1 9 5 4 Table 11.
377 dD ~d(H+)
Solubility of icetylacetone in Water as a Function of Ionic Strength at 25" C.
Grams
SaC1/25 MI.
Sq; Root of
Ionic Strengtli
Absorbance at 700 mp
0 0.094: 0.162
0.440 0.440 0 440
n
(1.0130 0 0384
0.287 0.410 0.644 0.750 0.096
0 1200
0.2451 0.6066 0 8264 1 4421
nK'(HR):-' (H+)"
d(HR), -- nK'(HR): ~d(H+) (H+)"+l
(5)
If--(l(HR)m
(HRw) the two right-hand trims of Equation 5 d(H+) - (H+) ' (HRL lircorne equal, but if is substantially smaller than __ (H+'i* ' dD nR'(HR):: d ( H again ~ becomes equal to and there should be ( H ')" + I no effect on the slope of the extraction curve. I n acetylacetone-
Solubility of Acetylacetone, G./RII. 0.173 0.173 0.173
0 432 0 423 0.398 0.383 0 346
-
,
0.166 0.159
0.146 0.139 0.123
\-
I
n-nter systems a t a p H of 0.5, (HR)wis approximately 50, but
(H +) ~ _ _ is much less than 1. Thus, even though acetylacetone
ronccntration of acid a t lower pH values. However, acetylawtoric is more soluble in a sulfuric acid solution than it is in :I sodium chloride solution having the same ionic strength. This is perhaps caused by the basic behavior of acetylacekone reprcsenteti by Equation 2 in which H R signifies the undissociated accty1:tc~etone.
+ Hf
HR
= HZR+
becomes less soluble in water at loa. pH values, there should lie little effect on the slope of the extraction curve, as contrasted with irstems in which
through distribution of
ilic reagent or through chelate formation. EFFECTSOF HYDROLYSISASD COMPLEXFORXATIOX. In :itidition to thr formation of thc chelate according to Equation 0
(2)
From the data of Tables I and I1 it is possible to calculate thc acid dissociation constant of the species (HzR)+. At a given ionic strength tht: difference between the solubility of acetylacetone i n sodium chloride solution and sulfuric acid solution wis assuInod to be the amount of the species (H?R)+, while the solubility in sodium chloride solution was assumed to represent thv xniount of the undissociated acetylacetone present. From thesis values t>heaverage acid dissociation of (H,R) a t several value.: of ionic strength was found to be 5.0 (Table 111). This valuc~ seems extraordinarily high, since it implies that acetylacetone is much more basic than similar compounds. Examining the (,onslants of benzoic acid determined bj- its increase in solu i n sulfuric acid solutions, Hammett ( 4 )says: "Obviously sulfuric, :icid may dissolve these substances by virtue of some other proljvrty t,han that of converting them to their conjugate acids."
-is large either
+ nR-
K
'
, IIR,,
7
(h)
the metal ion may be involved in coordination with other an1011.i. wch as chloride or the hvdroxyl ion, which may be repreqentcti I ) \ Equation 7
-
(7)
+
Calculation of the Ionization Constant of the Conjugate .4cid of icetylacetone H +. Square Root of IIR, HzR ,: Ionic Strength 3Ioles/Liter .\Ioles/Liter Mole/Liter K
Table 111. ,111 0
0.3 1 0 0 3
1.23 1.50 1.65 1.41
1
1 0.562 0.316 0.708
3.16 X 10-1 0 5.01 x 1 0 - 1
0.22 0.103 0.036 0.137
5.6 4.;
by.
5 0
4.1
5 2
Snvertheless the concentration of xetylacetone in water remains constant down to a p H of approximately 1. Therefore, there should be little difficulty i n obtaining extraction curves, niaintaining constant reagent concentration, (HR),, in the p H region 1.5 to 7 . Even below pII 1.3, the change of acetylacetone concentration with pH is much less than the change in a system rmploying an inert solvent. Between p H 2 and 0, the concentration of acetylstcetone in satur:tteti solutions changes from l . T
to 1.3moles per liter. Or
4H 1
=
0 4. It is interesting to see
what effect small changes of reagent concentration have on thP distrihution function. If the Kolthoff-Sandell extraction q u a tion
is differentiated keeping (HR', constant, Equation 4 results
dD d(H') ~
= - - .nk" (HR): ( H +),+I ~
If the term (HR), is a variable. differentiation gives:
(4)
In thiq
case
the following distribution function may be derive 1 :
+
nlielc 13 = 1 Kz(X--). .h pH increases I) ordiniirily increases. However, if t h e ( oncentration of X - increases-i.e., if 13 increases-with increabing pH as it would if X - were. for example, the hydroxyl ion, I ) increases less than it otherwise would. This means that tlic, resulting extraction curve would be less steep. Conversely, if the concentration of X - decreases with increasing pH, the slopc of the extraction curve will be steeper than it ordinarily Jvould be. If, ho~l-ever,the concentration of X- remains constant as pH varies, the shape of the curve will not change, since the term, p, drops out when D is differentiated with respect to ( H + ) In this case the curve will be translated without distortion toward high:r pH values. The extent of the shift or change of slope will depend on the magnitude of ( K z ) ( X - )as compared with 1. Since the sulfate complexes of the metals studied in this work possessed formation constants numerically less than 1 and since the sulfate concentration is never greater than 1, it was not expected that complexing should affect the extraction datn.
Table IT.
Distribution of Aluminum i n Presence of Fluoride
(Aluminum originally in acetylacetone phase at a concentration of 0.0154 mole/liter; fluoride concentration, 0.068 mole/liter; absorbance of original solution = 0.363) PH 6.90 4.78 2 47
Absorbance of hcetylacetone Layer 0.813 0.120 0 012
Amt. Extracted into Acetylacetone, % 50.2 33.1 3.3
JVhf>nan aluminum curve was run ~ 5 1 t hfluoride present in the aqueous laj-er, the extraction curve was far less steep than it wa* with no fluoride present (Table IV). This is in accord with the above considerations; since hvdrofluoric acid is a weak acid the fluoride ion concentration will increase as pH increases. Qualita-
ANALYTICAL CHEMISTRY
378
PcKiK?. . . K,t(HR)C(K,)" tive experiments involving acetylacetone (11) D = (H+)3 Ki[(HR)K,(H+)' K l K i . . . K,,1(HR):-'(K,)"-'(H+)'(11) extraction of iron(II1) solution containing large concentrations of phosphate indicate or for trivalent ions: that iron remains almost completely in the aqueous phase a t a p H of 2, whereas PcKiKzKa (HR):(K,)3 (12) (I1') it is normally completely extractedat this (")3 + K , (HR)K,(H+)z + K ~ (HR)Z K ~ ( ~ , ) z (H-) PH I n Figure 1 it can be seen that all three If the first term in the denJminator, (H-)3, is relatively much extraction curves have different slopes, although aluminum, gallarger, the second and third terms may be neglected and the prelium, and indium are all present as trivalent ions, which accordviously mentioned equation ( 5 , 6) for trivalent ions will remain ing to the Sandell-Irving equation should show equal slopes. in effect. However, if the third term is so large that the first and The curves may be compared to the theoretical monovalent and second terms may be neglected, the equation will reduce to the trivalent curves obtained by solving Equation 1 for D , using arbiform of the extraction equation of a monovalent ion. trary values for p H and K . It is noted that the aluminum curve From the ivork of Fernelius et al. (11) the minimum first and follows the monovalent curve. The gallium curve is slightly second formation constants for aluminum acetylacetonate in steeper, while the indium curve is almost parallel to the trivalent water may be estimated to be 1Oloand lo9, respectively. With slope. p H equal to 2, substitution of these values gives the following reThe variation in the slopes cannot be attributed to experispective values for the first, second, and third terms in the denommental error, for the maximum error in the determination of per inator: 1 X 10-6, 1.7 X cent extracted is about one or two units and the p H may be deterand 2.9 X Thus the second mined with an experimental error of f 0 . 0 3 p H unit. Furtherand third terms are nearly 100 times as large as the first, and on this basis i t is not surprising that aluminum has a lower slope more, the curves were reproducible. than anticipated from the simpler equation. (Since this manuSince both aluminum and gallium ions are relatively smaller script was written, Fernelius ( 2 ) has suggested that more recent than the indium ion and hence more subject to hydrolysis, this might be offered as an explanation of the lower slopes shown by work from his laboratory would make lo8 and 107 more likely aluminum and gallium. However, Bjerrum (1) gives the followvalues for K , and Kz, respectivelj-, of aluminum acetylacetonate. With these values, the values of the second and third terms in ing equilibrium constant for the hydrolysis of aluminum ion : the denominator become l . i X 10-5 and 2.9 X respectivelv, A1+3 HzO = AI(OH)+* H + still, it must be noted, larger than the first term.) However, if the formation constant for indium acetylacetonate is considerably lower than that of aluminum, the second and third terms become or smaller, and there should be greater conformance to the expected (.41(OH)+z) - (1.4 X trivalent slope. As shown in Figurr 1, the indium curve agrees (9) with the theoretical trivalent slope. Therefore, the authors feel (A1+3) (H+) that these additional terms should be considered in the extracFrom this it can be seen that the ratio of the hydrolyzed to tion processes involving metal chelates. The additional terms unhydrolyxed aluminum ions is very small in the region p H 1 to become negligible if the reagent concentration is very small. p H 3 in which the steepest portion of the extraction curves is Further, these terms are important only when the chelate has a located. Hence hydrolysis should have little effect on the slope rather high stability and also when the p H is relatively high. of the extraction curves. Substitution of approximate values in Equation 12 should indiEven though it is known (5') that aluminum forms complexes rate whether or not these terms are of significance. with such ions as sulfate. the decrease in slope cannot be attribThe position of the extraction curves along the p H axis deuted t o complex formation with the anions of the acid used to pends on the stability of the chelate extracted. Therefore aluacidify the water layers because the concentration of acid anion minum, having the highest chelate stability, should be extracted increases as p H decreases and this should result in a steeper rather a t the lov,est pH and gallium and indium at respectively higher than a less steep curve. Further, the same slope was obtained pH's. As shown in Figure 1 gallium i p extracted before indium, for aluminum with hydrochloric, nitric, and sulfuric acids embut aluminum is extracted a t an intermediate pH. The unexpectployed as acidifiers, and it is unlikely that the anions of all these edly high pHl,z value for aluminum is attributed to its relatively acids form complexes of equal stability with aluminum and thus low p , value which will shift the extraction toward higher pH's. equally affect the curve. Finally, a curve obtained with chloride .4s shown in the experimental part, iron may be removed from ion concentration varying with p H was identical with a curve in aluminum by extracting an aluminum solution a t a p H of 0.5 which the chloride ion concentration was maintained constant with acetylacetone. Preliminary experiments have shown that throughout the pH range. iron has a P H , , ~value of 0. Thus iron may be quantitatively Neither can the decrease in slope be ascribed to a change in removed from solutions of aluminum in lLV acid by extraction the concentration of acetylacetone in the water phases as the pH with acetylacetone. Hence, extraction curves for metals in decreases, since solubility determinations showed that the soluacetylacetone-water systems not only furnish interesting theobility of acetylacetone remains essentially constant a t 17.0 grams retical information, but may also indicate useful analytical separaper 100 ml. of solution down to a pH of 1. illso Sachod (8) has tions. demonstrated that the percentage of enol form of acetylacetone is essentially independent of pH in this range. ACKNOWLEDGMENT On reexamination of the extraction equations it is seen that the The authors are indebted to the U. S. Atomic Energy Comover-all formation constant, K,, is employed and the stepwise mission for financial support. formation constants following were not considered.
+
+
+
&f+"
MR+("-')
+
+ R-
+ R-
=
&JR+(n-C
K,
MRz+("-*)K,, etc.
LITERATURE CITED
(10)
If these equations and the resulting equilibrium constants are employed in the derivation of the distribution function Equation 11 obtains:
(1) Bjerrum, K., 2. p h y s i k . Chem., 59, 350 (1907). (2) Fernelius, U '. C., private communication. (3) Guiter, H., Compt. rend., 226, 1092 (1948). (4) Hammett, L. P., "Physical Organic Chemistry," York, McGraw-Hill Book Co., 1940.
p. 273, New
V O L U M E 2 6 , NO. 2, F E B R U A R Y 1 9 5 4 ( 5 ) Irving, H., and Williams, J., J . Chem. Soc., 1949, 1841. (6) Kolthoff, I. AI., and Sandell, E. B., J . Am. Chem. Soc., 63,1906
(1941). (7) Morgan, G. T., and Drew, H. D. K., J . Chem. Soc., 1921, 1058. (8) Kachod, F. C., 2. p h y s i k . Chem., A182, 193 (1938). (9) Siggia, S., “Quantitative Organic Analysis via Functional Groups,” p. 111, ?;em- York, John Wiley &- Sons, 1949.
379 (10) Steinbach, J., and Freiser, H., SNAL. CHEM.,25, 881 (1953).
(11) Van Gitert, L. G., Fernelius, W. C., and Douglas, B. E., U. 9. Atomic Energy Commission, R e p t . , NYO-3372. RECEIVED for review J u n e 26, 1953. Accepted Norember 19, 1953. Presented before t h e Division of Analytical Chemistry a t the 124th Meeting of the AMERICAN CHEMICAL SOCIETY,Chicago, Ill., 1953. Contribution No. 918 from the Chemistry Department, University of Pittsburgh.
Direct Photometric Determination of Iron, Titanium, and Copper in Tantalum and Its Oxide JANE HASTINGS, T. A. MCCLARITY, and E. J. BRODERICK Materials and Processes Laboratory, Transformer Division, General Electric
The work described was undertaken in order to provide rapid chemical methods for the analysis of pure tantalum metal and to furnish data for the calibration curves for spectrographic examination of the material. The spectrophotometric determination of iron by ophenanthroline was used in the range of 0.002 to 0.030qo of iron, of copper by sodium diethyldithiocarbamate in the range of 0.000 to 0,01670 of copper, and of titanium by hydrogen peroxide and sulfuric acid in the range of 0.004 to 0.070qo of titanium. Calibration data are for each element as obtained with a Beckman DU spectrophotometer. Tedious separations of the elements to be determined are unnecessary and contamination from the use of large quantities of reagents is avoided. A tenfold gain in time was realized for the analysis of samples of tantalum metal. Precision and accuracy have been greatly improved.
DETERXIINATIOY OF IROY
General Principles. The tantalum metal is dissolved in hydrofluoric and nitric acids. Most of the acids are volatilized and a tartaric acid solution of the salts is made a t once. The iron is reduced in the presence of excess tartaric acid by hydroxylamine hydrochloride solution. o-Phenanthroline solution is added and the formation of the red ferrous-o-phenanthroline is the basis for the determination of iron ( 5 ) . There are no interferences to be expected in commercially pure tantalum ( 3 ) . 3 60C A C u r v e fa solution without t a n t a l u m
0 Curve l o r solution containing t a n t a l u m i 1 0 0 0 g ) A = 5 0 8 m p I c m cells Slit w i d t h =O 2 6 m m D i s t i l l e d woter: 100%transmission 0 50C
THE
growing use of pure, and in some cases, comparatively unfamiliar metals and the necessity for evaluating the elements present in small amounts in them, have forced the analytical chemist to look for means of extending the sensitivity of The chemistry of tantalum is complex and the problem of analyzing the minerals in a hich the element occurs as the oxide is covered classically by the work of Schoeller ( 6 ) . The separations and determinations are tedious and time-consuming and do not yield consistent results as shown by Atkinson et al. (1). The methods presented in this paper were developed to meet the need for analyzing a large number of samples of tantalum metal for minor constituents. Lengthy separations and contamination of the end products of the analysis by the necessity for using large amounts of reagents are avoided. Although standards were not mailable, by using a single sheet of tantalum metal as a reference for all calibration curves, it has been possible to evaluate the reproducibility of the method. The accuracy of the methods is believed to be good, too, since the results obtained agreed well T+iththose obtained when iron, titanium, and copper were separated b r conventional methods and determined photometrically. ill1 of the procedures take advantage of the fundamental reactions that tantalum and niobium form soluble complexes with tartrates and oxalates as developed by Schoeller ( 6 ) and his associates. The solutions are stable enough to provide media for colorimetric reactions which are specific for the elements to be determined. The danger of hydrolysis, characteristic of tantalum salts in solution, is still present but proper attention to detail will yield reproducible results. The Beckman DU spectrophotometer was used for all measurements of rolor.
Co., Pittsfield, Mass.
0 40(
> 0 z
s
030C
---
a
G
2 0 20(
0 IO(
0 OO(
Figure 1.
Calibration Curves for Ferrous o-Phenanthroline Complex
The method has been found to be reproducible through the range of 0.002 to 0.030q70 of iron. Blanks on reagents carried through the procedure showed an average of 0.002y0 iron content with a spread of 0.0004% from day to day. The Lambert-Beer law, as shown in Figure 1, is obeyed for the solutions and the presence of tantalum does not affect the determination seriously.
.
