W is the weight of sample S is the sulfate content of the samale a s % NazSO,
If the normality of the sodium hydroxide solution is not exactly 0.1, t~heabove factors should be changed according t o the true normality.
For cryolite 0.95(T
-0
w
05) -
268s
1.75(T - 0 05)
F%
-___
w
=
ANALYTICAL RESULTS
- 0.4939 =
Na3.41F6%
For fluorspar 0.76(T - 0.05)
w
1.562(T - 0.05) W
=
=
F%
CaF2%
where 0 . 0 5 represents a column and indicator blank T is the net titration
Tables I, 11, and I11 show typical results for aluminum fluoride, cryolite, and fluorspar, respectively. I n all cases, the values shown represent the total fluorine content calculated to the compound. I n these tables, distillation method refers to the steam distillationthorium nitrate titration procedure. The lead chlorofluoride method referred to in Table I is that recommended by Specht and Hornig (7).
(2) Berl, E., Lunge, G., “Chemischetechnische Untersuchungsmethoden,” 8th ed., Vol. 2, Part 1, p. 279, Julius Springer, Berlin, 1932. (3) Hawley, F. G., Ind. Eng. Chem. 18, 573 (1926). (4) Kolthoff, I. hl., Stenger, V. A , “Volumetric Analysis,” 2nd rev. ed., Vol. 2, p. 121, Interscience, New York, 1947. ( 5 ) Proske, O., Blumenthal, H., “Analyse der Metalle,” 2nd ed., Vol. 1, p. 19, Julius Springer, Berlin, 1949. (6) Schucht, L., Moller, W., Ber. 39, 3693 (1906). (7) Specht, F., Hornig, A., Z. anal. Chem. 125, 161 (1943). ( 8 ) Spielhaczek, H., Ibid., 118, 161 f193R) (9) Treadwell, F. P., Hall, W. T., \ - - - - I .
“Analytical Chemistry,” 9th ed., Vol. 2, p. 398, Wiley, New York, 1942. (10) Willard, H. H., Winter, 0. B., IND. E N G . CHEM., ANAL. ED. 5 , 7 (1933).
LITERATURE CITED
(1) ildolph, W. J., J. Am. Chem. Sac. 37, 2500 (1915).
RECEIVEDfor review February 7, 1957. -4ccepted May 13, 1957.
Precipitation of Magnesium with (Et hylenedinitri Io)tetra acetic Acid CLARK E. BRICKER and GREGORY H. PARKER’ Frick Chemical laboratory, Princeton University, Princeton, N. J.
b The use of (ethylenedinitri1o)tetraacetic acid (EDTA) as a complexing agent is well known but the fact that it can b e used as a reagent for the precipitation of magnesium has not been reported. Between pH 3.5 and 4.0, magnesium is precipitated slowly and quantitatively in the presence of an excess of the disodium salt of EDTA as a salt which closely approximates MgCloH14O8N2.6H20. The precipitate can b e dried a t room temperature and weighed as the hydrate, or vacuum-dried a t 100’ C. and weighed as the anhydrous salt. Either weighing form has a very favorable gravimetric factor for the determination of magnesium. Interference studies with 13 cations and six anions show that EDTA is not a specific reagent but is selective as a precipitant for magnesium. The properties and the rate of precipitation of this magnesium salt were studied.
s
(ethylenedinitri1o)tetraacetic acid (EDTA) was first made available by the I. G. Farbenindustrie in 1936 (4), the use of its disodium INCE
* Present address, Electro Metallurgical Co., 137 47th St., Nisgara Falls, N. Y. 1470 * ANALYTICAL CHEMISTRY
salt as a complexing agent has found numerous applications. Pecsok (6) gives ten general categories for the analytical use of the sodium salts of EDTA, but all of these applications depend in some way on the formation of a soluble complex between this reagent and the ions of the metal(s). Brintzinger and associates (1-3) reported the preparation of heavy metal compounds of EDTA from highly concentrated solutions, producing both simple salts, with the metal bonds attached only to the carboxyl groups as in the disodium salt, and inner complex salts with the typical chelate structure. The reported solubility of these salts varied widely but was generally high. Lanthanum, on the other hand, formed a very sparingly soluble salt and although Brintzinger suggested the possibility of analytical application, no such method has been reported. This paper deals with the precipitation of magnesium dihydrogen (ethylenedinitri1o)tetraacetate and the use of this precipitate for separating and determining magnesium. EXPERIMENTAL
During an investigation of the possibility of separating magnesium and
calcium on an ion exchange resin by selectively complexing the calcium with EDTA, a white crystalline material was observed to precipitate slowly from a dilute solution of magnesium and EDTA ions at p H 3.50. When the solution was made alkaline with sodium hydroxide, the precipitate did not dissolve on standing overnight. Therefore, it was concluded that the precipitate was not EDTA, as this substance would be readily soluble in alkaline solution. As a solution of calcium and EDTA ions produced no precipitate under the same conditions, a series of investigations was undertaken to determine the formula and properties of the precipitate. Apparatus and Reagents. A Beckman Model D U spectrophotometer equipped with a Model 4030 atomizerburner and a Model 4300 photomultiplier accessory was used for the flame photometric measurements. A Fisher Electrophotelometer with a 5 2 5 - m ~filter was used for the photometric determinations of magnesium with Titan Yellosy ( 5 ) . All pH measurements were made on a Leeds &- Northrup KO.7664 p H meter, using a KO. 1199-30 glass electrode and a KO. 1199-31 reference electrode. The stirring motors were heavy-
duty, variable-speed motors and were operated a t high speed. The Vibromixer was supplied by the Fisher Scientific Co. (Cat. Xo. 14-510-50). A weighed amount of disodium dihydrogen (ethylencdinitri1o)tetraacetate was dissolved in water and the resulting solution was standardized against a standard calcium solution by direct titration to an Eriochrome Black T end point and found to be 0.0948iM. Several other solutions of this reagent were prepared by weight, so that they were approximately 0.1M. A magnesium solution was prepared by dissolving reagent grade magnesium oxide in hydrochloric acid, evaporating, and diluting to volume. This solution 11 BS standardized against the standard EDTA solution and found to be 0.05S9M. Another magnesium solution !%asprepared from pure magnesium ribbon and was 0.1004M. The standard calcium solution was prepared by dissolving a weighed amount of dried and cooled calcium carbonate (analytical reagent grade) in water plus the minimum amount of 6N hydrochloric acid. The resulting solution was diluted to volume in a volumetric flask. Solutions of the various metallic ions that were used for interference studies were prepared from reagent grade chemicals and varied from 0.10 to 0.50M. Formula of Precipitate. When the precipitate obtained from knowii amounts of magnesium ions and disodium dih ydrogen (et hylenedinitrilo) tetraacetate was dried a t room temperature and weighed, 17.4 mg. of precipitate were found for each milligram of magnesium in the precipitate. The filtrate from one magnesium EDTA precipitation was titrated with standard magnesium solution to an Eriochrome Black T end point to determine the amount of excess EDTA present. The results of this experiment proved conclusively that there was a ratio of one magnesium to one (ethylenedinitriloftetraacetate in the precipitate. The molecular weight of the magnesium compound is, therefore, calculated to be 423. When some of the precipitate was dissolved in water and the solution was analyzed with a flame photometer for sodium, the test was negative. The formula, h'IgC10H1408r\T26H20, with a molecular weight of 422.6, was proposed for the precipitate. When samples of the precipitate, which were dried a t room temperature in a desiccator, were analyzed for water by the Karl Fischer method, only 4.60 to 7.65% of water was found as compared t o 25.58% of water in the proposed formula. These low results Jvere attributed t o the slight solubility of the precipitate in absolute methanol and to the tenacity with which the water is apparently held in the molecule. However, when 232.4- and 239.5-
mg. samples of the precipitate were vacuum dried a t 98" C. to constant weight, 59.4 and 61.2 mg., respectively, were lost. These losses correspond to 25.56 and 25.55y0of water. Samples of the magnesium EDTA precipitate were ignited to constant weight to magnesium oxide in porcelain crucibles (Table I). Samples 1 and 2 had been dried a t room temperature and, on the basis of the proposed hexahydrate should have contained 9.547, of magnesium oxide by weight. Sample 3 was vacuum-dried a t 98" C. and, if the anhydrous salt, it should have contained 12.82y0of magnesium oxide. Therefore, because the precipitate of magnesium EDTA contains no sodium, equimolar amounts of magnesium and EDTA, 25.6% of w t e r , 9.7% of magnesium oxide, and a calculated molecular weight of 423, the material must be magnesium dihydrogen (ethylenedinitrilo) tetraacetate hexahydrate. Factors Influencing Precipitation. To pursue the analytical applications of the precipitate, it was necessary to determine optimum conditions for the precipitation Ivith completeness of removal of magnesium from solution as the criterion. The general procedure consisted of pipetting the desired amount of the standard magnesium and EDTA solution into 100-ml. beakers and adding sufficient distilled ivater to give 35 to 40 ml. After the pH of the solution was adjusted to the desired value with 2iV sodium hydroxide, the beakers were
Table I.
Sample 1
2
3 Table II.
covered with watch glasses and allowed to stand the required length of time. The precipitates were washed from the beakers t o weighed, medium-porosity, fritted crucibles with a rubber policeman and a jet of water, then washed with approximately 5 ml. of water, and finally with 5 ml. of acetone. The crucibles were then placed in vacuum desiccators, with anhydrous magnesium perchlorate as the desiccant, and dried to constant weight a t room temperature. The magnesium EDTA precipitate corresponds closely to the hexahydrate under these conditions and reaches constant weight in about 5 hours. Generally, no appreciable amount of precipitate could be observed for 3 t o 5 hours after mixing and adjusting the pH of the magnesium and EDTA solutions. After about 5 hours a small amount of precipitate was visible on the bottom of the beaker, and the number and size of these particles could be observed to grow in size for about 1 day. The precipitate was generally very granular and very easily removed from the beaker. However, when high concentrations of magnesium Ivere precipitated, the magnesium EDTA precipitated more rapidly in a finely divided form and underwent no observable change after aging a t room temperature. pH. The optimum pH for the precipitation was immediately established between 2.0 and 5.0, since below pH 2.0, EDTA precipitates and above pH 5.0, magnesium EDTA precipitates
Ignition of Magnesium EDTA to Magnesium Oxide
Sample Taken, Mg. 399.9 197.6 92.5
Wt. of Ignited Sample, Mg. 38 5 19.3 11.5
Theoretical Wt. of Ignited Sample, Mg. 38 2 18 9 11.9
bIgO Found, % 9.63 9.77 12 4
Recovery of Magnesium as a Function of EDTA Concentration and Time of Standing before Filtration
EDTA soln. Taken, 0.1M, M1. 25 20
Time of Initial Standing, Wt. of Ppt., Mga pH Hours Mg. Recovery, % 3.70 48 240.2 96.7 3.70 48 238.8 96.0 15 3.70 48 231 .O 93.0 10 3.70 48 147.4 59.4 25 4.00 72 245.9 98.8 20 4.11 72 239.2 96.0 15 4.00 72 237.1 95.4 10 4.00 72 232.1 93.4 3.70 12 days 30 247.2 99.4 25 3.70 12 days 247.9 99.7 20 99.7 3.70 12 davs 247.9 15 90.2 40 3.70 12 dairs 246.5 10 12 days 245.4 40 3.70 98.7 50 51 3.79 96 30.1 71.0 25 1.4 3.3 51 3.83 96 10 0.0 0.0 51 3.81 96 a 14.3 mg. of magnesium present in all cases except in the last three results, when only 2.44 mg. of magnesium were added. Per cent recovery of magnesium was based on assumption that precipitate was MgCloHlrO8N2.6H20. Final Vol., M1. 40 40 40 40 35 35 35 35 40 40 40
~
VOL. 29, NO. 10, OCTOBER 1957
1471
extremely slowly, if a t all. The optimum pH for the precipitation was established through the use of buffered solutions to be betm-een 3.5 and 4.0. Because the pH of unbuffered solutions of magnesium and EDTA rises 0.2 t o 0.3 pH unit during precipitation, the optimum pH for starting the precipitation should be about 3.70. Concentration and Time of Standing. The results of a number of experiments in which the concentration of the E D T A and the time of standing were varied are shown in Table 11. It is obvious from these results that a better recovery of magnesium is realized the longer the solution stands prior to filtration. The highest recoveries for 14.3 mg. of magnesium were obtained when a molar ratio of EDTA to magnesium of a t least 4 to 1 was present. The last three results in Table I1 are particularly interesting because small amounts of magnesium do not precipitate quantitatively even though the molar ratio of EDTA to magnesium is as high as 50 to 1. Because the concentration of reactants and time of standing prior to filtration are interrelated for the quantitative precipitation of magnesium EDTA, it is difficult t o state any optimum conditions. However, if a solution is allowed to stand quietly a t room temperature 48 to 72 hours prior to filtration, the optimum concentration of reactants for the quantitative precipitation of magnesium EDTA is given by the empirical expression:
V = 13 ( A where B
=
A
=
V
=
+B)
millimoles of EDTA present and is a t least 4 A millimoles of magnesium present total volume of solution in milliliters
Stirring. Using the optimum p H and optimum concentration of reactants, the authors found t h a t the rate of precipitation of magnesium E D T A was increased approximately 30 times by stirring the solutions rapidly with a motor-driven glass rod with a horizontal disk a t the end. Extreme agitation with a Vibro-mixer appeared to increase the rate of precipitation even more but, because of the difficulty of keeping the solution quantitative with this type of stirring, only qualitative experiments were conducted with this apparatus. Even though the magnesium EDTA is much more finely divided when precipitated from stirred solutions, its composition is identical \vith that obtained from unstirred solutions. The finely divided magnesium EDTA, after it is formed, does not appear to increase in crystal size on aging in quiet solution. This is probably due to the slow rate 1472
ANALYTICAL CHEMISTRY
Table 111.
Recovery of Magnesium from Mixed Solvent Systems pvlg
Solvent Isopropyl alcohol Ethyl alcohol Dioxane Table IV.
Sample Taken, Mg.
282.8 85.6
"g
Taken,
W-t. of Ppt.,
Mg.
Mg.
in Filtrate, Mg.
14.3 14.3 14.3
260.0 254.5 307.0
0.4 0.1 0.0
wt. of Ppt., per Mg. of
Magnesium 18.7 17.9 21.4
Equilibrium Solubility of Magnesium EDTA in Water
Loss in
Sample Recovered,
Weight,
Mg.
Mg. 14.6 15.6
268.2 70.0
a t which the precipitate dissolves in water a t room temperature. Mixed Solvents. I n order to determine whether a mixed solvent system would increase the rate of precipitation of magnesium E D T A , solutions of magnesium and EDTA were prepared a t pH 3.70 containing approximately 42% by volume of the nonaqueous solvent, The total volume of each solution was 60 ml., which n-as 50y0 greater than the optimum volume in aqueous solution for the amounts of magnesium and EDTA taken. The solutions were allowed to stand quietly for 24 hours a t room temperature and were then filtered, and the precipitates were vacuum-dried a t room temperature. The filtrates from these experiments were analyzed spectrophotometrically for magnesium (6) (Table 111). As contrasted to all aqueous solutions, the solutions of mixed solvents containing magnesium and EDTA appeared turbid almost immediately and a definite precipitate was visible in 15 to 20 minutes. These precipitates, however, appeared somewhat flocculent and did not settle rapidly. A portion of the precipitate from the dioxane system was dissolved in water and this solution gave a strong test for sodium on the flame photometer. Vacuum oven-drying of this precipitate at 98' C. and ignition of the residue t o magnesium oxide and sodium carbonate gave results which suggested no simple compound but rather a mixture of magnesium EDTA and possibly disodium dihydrogen (ethylenedinitri1o)tetraacetate dihydrate. The fact that magnesium is quantitatively precipitated in 24 hours in the dioxane system may suggest a rather selective method for separating magnesium. However, the uncertain composition of the precipitate precludes the use of this method for the gravimetric determination of magnesium. PROPERTIES OF MAGNESIUM EDTA
Solubility. I n order t o determine the solubility of magnesium E D T A in
Solubility, Mg./MI. HX0 0.58 0.62
rTater, two samples of widely different weights were placed in separate 25.00ml. portions of water, refluxed for 9 hours a t a slow boil, allowed to recrystallize for 7 days at room temperature, filtered, dried, and weighed. The smaller sample dissolved completely during the reflux and the recrystallized material consisted of larger crystals than the initial sample. The data from these experiments (Table IV) show the solubility to be independent of the sample size and to have an equilibrium value of 0.60 mg. of magnesium EDTA per ml. of water a t room temperature. Solubility loss in washing precipitates of magnesium EDTA with water and acetone was found t o be negligibleLe., less than 0.05 mg. of magnesium EDTA when 10 ml. of water and 10 ml. of acetone were in contact with the precipitate for not more than 5 minutes. This is true because the precipitate dissolves very slo~-lyin water and, therefore, the equilibrium solubility is never approached by the wash water. Vapor Pressure of MgCI0Hl4O8N2.6Hz0.A sample of desiccator-dried magnesium E D T A was placed in a sealed, evacuated, all-glass apparatus and the pressure in this system was measured by a closed-end mercury manometer. The entire apparatus was immersed in a constant temperature bath. The temperature, time allowed to reach equilibrium, and the vapor pressure a t each temperature are given in Table V. These data indicate that the magnesium EDTA hexahydrate does not show any appreciable vapor pressure until a temperature of at least 40" C. is reached, Furthermore, the hydrate is stable at elevated temperatures and verifies the need for vacuum drying the precipitate a t 98" to 100" C. in order to obtain the anhydrous salt. The vapor pressure of magnesium EDTA does not appear t o be a reversible phenomenon, because the vapor pressure of the heated salt did not return to its original value when the temperature was reduced to 25" C. (last line in Table V). The anhydrous
salt was not noticeably deliquescent in 50 to 70% relative humidity and, therefore, no difficulty was encountered in weighing the dried salt. RECOMMENDED PROCEDURE FOR DETERMINATION OF MAGNESIUM
To the solution containing from 0.2 to 1.0 mmole of magnesium in 10 to 15 ml., add 25 ml. of O . l M disodium dihydrogen (ethylenedinitri1o)tetraacetate and adjust the p H of the solution with either sodium hydroxide or hydrochloric acid to a value between 3.6 and 3.8. Allow the solution to stand for approximately 48 hours a t room temperature and then filter through a weighed fritted-glass crucible. Wash the precipitate with 10 ml. of water and 10 ml. of acetone and then dry the precipitate in a vacuum desiccator over magnesium perchlorate to constant weight. Because the precipitate obtained in this procedure contains on the average 24.5% of water instead of the calculated 25.6% of water for bIgCloH1108N2.6H20, a n empirical factor of 0.05832 instead of the theoretical 0.05754 should be used to calculate the weight of magnesium in the precipitate. On the other hand, if the washed precipitate is vacuum-dried in a n oven at 98" t o 100" C. for approximately 24 hours, the anhydrous salt is then weighed and the theoretical factor of 0.0772 is used to calculate the weight of magnesium in the precipitate. If the solution contains, in addition to magnesium, over 1.0 mmole of cations that can form soluble complexes with E D T A a t a p H of 3.5 to 4.0, it is necessary to add more reagent to complex these ions. This additional precipitant is needed especially if appreciable amounts of the cations from the elements of the first transition group or trivalent or quadrivalent cations are present but not with calcium, barium, beryllium, lithium, or sodium. RESULTS AND INTERFERENCES
The results of some typical analyses by the recommended procedure either with or without possible interfering ions present are given in Table VI. The results for the solutions containing only magnesium show a n average recovery of 99.96% of magnesium based on the weight of the hydrated precipitate and of 99,7770 based on the weight of the anhydrous salt. Barium and lithium do not interfere appreciably with the separation and determination of magnesium. The interference of calcium is small when the weight of this ion present is less than that of the magnesium. For larger amounts of calcium the interference is significant but can be virtually eliminated by a double precipitation (footnotea, Table VI). Beryllium seems t o interfere somewhat by causing the precipitate of magnesium E D T A t o have less de-
Table V.
Vapor Pressure of
MgCioHirOeNz.6HzO Time at Manometer Temp., Reading, Temperature, O C. Hours Mm. Hg 30 ... 0.0 40 12 0.2 49.5 26 2.0 55 12 2.2 60 18 3.0 85 6 8.5 25 4Days 3.5
sirable properties for filtering and washing. I n addition to the results shown in Table VI, the possible interferences of nine other elements were studied. When as much as 2.5 mmoles of cobalt(I1) , copper(I1) , nickel(II), zinc(II), cadmium(II), or thorium(1V) or 1.25 mmoles of bismuth(II1) were mixed with 0.60 mmole of magnesium ions, the recovery of magnesium, based on the weight of the hydrated precipitate, was over 99% in every case. Further-
Table VI.
more, none of these foreign ions was detected in the precipitates when they were tested by appropriate color reactions. On the other hand, with 2.5 mmoles of aluminum(II1) present, the recovery of magnesium was only 81%; with 2.5 mmoles of manganese(II), the recovery was 35700. A few experiments were carried out to determine the effect of acetate, borate, phosphate, tartrate, and citrate on the recovery of magnesium by the recommended procedure. With the precipitating solution, 0.1M in either borate and/or acetate, the recovery of magnesium was about 98%; with 0.1M phosphate, 93y0; with 0.1M tartrate or citrate, 89%. The possibility of reversing the recommended procedure and using excess magnesium ions to precipitate (ethylenedinitri1o)tetraacetate was studied briefly. Known amounts of (ethylenedinitrilo) tetraacetic acid were mixed with an excess of magnesium ions; the p H of the solution was adjusted to 3.7, and after the solution was allowed
Analyses by Recommended Procedure
Mg Interference ivt. of Wt. of Taken, Weight, Hydrated %a Anhydrous % Mg. Ion mg. Ppt. Recovery Ppt. Recovery 12.20 ... .. . 0.2074 99.14 0.1572 99.62 100,89* 99.86 0.1592 12.20 ... .. . 0.2089 12.20 ... .. . 0.2096 100.19 0.1576 99.85 12.20 ... ... 0.2089 99.86 0.1575 99.78 0.3351 99.82 24.41 . ., ... 0.4198 100.33 0.1576 99.85 12.20 ... ... 0.2094 100.10 12.20 ... ... 0.2097 100.24 0.15i4 99.72 12.20 Ba 27.0 0.2090 99.90 0.1586 100.5W 0,1589 100.70c 12.20 Ba 68.0 0.2089 99.86 0.1578 1OO.OW 12.20 Ba 137.0 0.2086 99.71 12.20 Li 1.4 0.2090 99.90 0.1576 99.87 12.20 0.2096 100.19 0.1580 100.13 Li 3.5 99.4gd 12.20 Li 7.0 99.62 0.1570 0.2084 12.20 Ca 3.8 0.2062 98.57 12.20 Ca 7.5 0.2089 99.86 ... . . . 0.2105 100.62 ... ... 12.20 Ca 11.3 12.20 Ca 18.8 0.2145 102.53 Ca 18.8 0.2162 103.34 o.'i& io4:is 12.20 4.88* Ca 37.6 0.0803 96.1 0,0614 97.3 12.20e Ca 37.6 0.2081 99.47 0.1568 99.24 103.86 103.39 0.1639 12.20 Be 1.8 0.2163' 0.1557 98.67 12.20 Be 4.5 0.2053' 98.61 a Empirical factor of 0.0583 used to calculate weight of magnesium in hydrated pre= cipitate. b Not counted in calculating average % recovery. c Spectroscopic analysis showed approximately 0.2y0 of Ba from combined precipitates. d Spectroscopic analysis showed 0.005 t o 0.01yoof Li. After following recommended procedure and filtering supernatant liquid through weighed fritted crucibles, 20 ml. of 0.05M disodium dihydrogen (ethy1enedinitrilo)tetraacetate Fere added to each precipit'ate. Resulting mixture was refluxed for 3 hours and allowed to stand 3 days before filtering, washing, drying, and weighing precipitat,e. Precipitate stuck to beaker more than other precipitates; crystals were plates. 6
0
Table VII. Recovery of EDTA from Aqueous Solution Wt. of Ppt., Recovery of EDTA Taken, Magnesium Added, Grams EDTA, %" Mmoles Mmoles 2.42 1.47 0.6148 60.0 2.42 2.95 1 ,0032 98.2 2.42 5.89 1.0763 105.0 Recovery based on assumption that precipitate is MgCloHl408N2.6HsO.
VOL. 29, NO. 10, OCTOBER 1957
1473
to stand 2 days, the precipitate was filtered, washed, and dried in a vacuum desiccator. The results, shown in Table VII, suggest that this procedure is adequate to remove (ethylenedinitri1o)tetraacetate from solution but not sufficiently accurate for a quantitative determination. A few experiments with Versene-ol and Versene-diol (obtained from the Dow Chemical Co.) showed that neither compound would precipitate magnesium under the same conditions as did (ethylenedinitri1o)tetraacetate. Khen
Versene-diol was mixed with (ethylenedinitri1o)tetraacetic acid, the rate of precipitation of magnesium EDTA was retarded and only 52y0 of the magnesium was precipitated. A mixture of Versene-ol and (ethylenedinitri1o)tetraacetic acid gave no precipitate a t all with magnesium. LITERATURE CITED
(1) Brintzinger, H., Hesse, G., 2. anorg. u. allgem. Chem. 249, 113 (1942). (2) Brintzinger, H., Munkelt, S., 2. anorg. Chem. 256, 65 (1948).
(3) Brintainger, H., Thiele, H., Muller, U., 2. anorg. u. allgem. Chem. 251, 285 (1943). (4) I. G. Farbenindustrie, Ger. Patent 638,071 (1936). ( 5 ) Ludwig, E. E., Johnson, C. R., IND. ENG.CHEY.. ANAL. ED. 14, 895 (1942). (6) Pecsok, R. L., J . Chem. Educ. 29, 597 (1952). RECEIVEDfor review March 2, 1967. Scce ted May 2, 1957. Presented before the Eighth Pitt.vburgh Conference on Analytical Chemistry and Applied Spectroscopy, Pittsburgh, Pa., 1957.
Separation and Determination of Tantalum GLENN R. WATERBURY and CLARK E. BRICKER' The University of California, 10s Alamos Scientific laboratory, 10s Alamos, N. M. ,Between 0.025 and 3.00 mg. of tantalum are extracted into hexone (4-methyl-2-pentanone) from 6M sulfuric acid-0.4M hydrofluoric acid media and estimated colorimetrically using hydroquinone in concentrated sulfuric acid. Tantalum found averaged 100.2%, with a standard deviation of 1.3%, in 50 analyses of known solutions containing 0.5 to 3.39 mg. of tantalum and various amounts of other elements. No interference in determination of 1 mg. of tantalum was caused by 100 mg. of sodium, potassium, uranium, and plutonium, 20 to 40 mg. of aluminum, chromium(Ill), strontium, vanadium(V), thallium(I), cobalt, and iron(lll), and 2 to 3 mg. of tungsten, neodymium, cerium, zirconium, ruthenium(lV), and gold. Interference of titanium and molybdenum may b e eliminated b y double extraction. Niobium is the only metal that interferes seriously.
B
an analytical procedure was needed for estimating 0.01 t o 2% tantalum in uranium and plutonium alloys, possible methods for the separation and determination of small amounts of tantalum were investigated. Unless large samples are taken, the usual gravimetric method for tantalum is not sufficiently accurate or sensitive. KO specific titrimetric method exists for the determination of tantalum. Of the general separation techniques, extraction methods have been investigated thoroughly and seemed to offer the most promise. Stevenson and Hicks (O),using radioECAUSE
1 Present address, Chemistry Department, Princeton University, Princeton, N. J.
1474
ANALYTICAL CHEMISTRY
tracer methods, found that tantalum could be extracted with diisopropyl ketone from a solution of tantalum and niobium that was 6M in hydrochloric or sulfuric acid and 0.4M in hydrofluoric acid. Other extraction techniques include removal of tantalum from 5M sulfuric acid with 8% tribenzylamine in methylene chloride ( 2 ) , separation of tantalum from 21V sulfuric acid using 5% methyloctylamine in xylene (6), extraction of the tantalum fluoride complex with ethyl acetate (11), and separation of tantalum and niobium from hydrochloric-hydrofluoric acid solution using methyl isobutyl ketone (7, I S ) . For determination of small amounts of tantalum after separation, colorimetric methods seemed to offer the sensitivity required. The pyrogallol (4, 6, IO), perhydrol (8), and chromotropic acid reactions (I) have been adapted for colorimetric procedures. By combining separation of tantalum by extraction with colorimetric determination, a satisfactory procedure has been developed. Although it is written specifically for analysis of samples of uranium or plutonium alloys, the general method, except for the dissolution steps, may be applied to any niobium-free sample that is soluble in 6M sulfuric acid-0.4M hydrofluoric acid. APPARATUS A N D REAGENTS
Platinum dishes, with lip. Hollow-tube type extractor, with 25 x 150 mm. test tubes, as shown in Figure 1. Flasks, 10-ml. volumetric, borosilicate glass. with glass stoppers. Infrared heat lamp with socket, switch, extension cord, and ring stand attachments.
Electric hot plate. Polystyrene pipet, 2-ml., with 0.5-ml. graduations. Beckman Model DU spectrophotometer, with Corex cells of 1-cm. light path. Hydrofluoric acid, 48%, reagent grade. Hydrofluoric acid, 4M. Dilute 16.26 grams of 48y0hydrofluoric acid to 100 ml. with water. Store in a polyethylene bottle. Nitric acid, concentrated, reagent grade. Sulfuric acid, concentrated, reagent grade. Dispense from a 25-ml. reservoir buret equipped with silica gel drying tubes on the air inlet and top vent. Ammonium persulfate crystals. Hexone (4-methyl-2-pentanone), Eastman white label, or methyl isobutyl ketone, Matheson, Coleman, and Bell. Hydroquinone solution, 55 mg. per ml. Dissolve 5.50 grams of Eastman white label hydroquinone in concentrated sulfuric acid and dilute to 100 ml. with this acid. Dispense from a 5-ml. Koch microburet equipped with silica gel drying tubes at the top of the buret and a t the reservoir vent. Sodium hydroxide solution, 5 N . Dissolve 20 grams of reagent grade sodium hydroxide pellets in water and dilute to 100 ml. Store in a polyethylene bottle. Standard tantalum solution, 0.6 mg. of tantalum per gram. Accurately weigh about 0.1 gram of pure tantalum metal into a 50-ml. platinum dish, and add 4 ml. of water and 2 ml. of nitric acid. Add hydrofluoric acid dropwise to maintain a slow dissolution rate until the tantalum is dissolved. Cautiously add 35 ml. of concentrated sulfuric acid and evaporate the solution to strong fumes of sulfur trioxide. Transfer the solution to a weighed 125-ml. glassstoppered flask and dilute to about 100 ml., using concentrated sulfuric acid to wash the dish and for the dilution. Weigh the flask and contents and calcu-
