Use of a Lead Reductor in the Potentiometric Determination of Uranium and Iron W. R. AMOS and W. B. BROWN Mound Laboratory, Monsanto Research Corp., Miamisburg, Ohio
b Iron and uranium were determined simultaneously b y potentiometric titration with potassium dichromate. The samples were passed through a lead reductor into sufficient potassium dichromate to oxidize 95y0 of the estimated uranium in the sample. After the resulting solution was purged with argon and heated to 85' to 95" C., the remaining 5% of the uranium(lV) was titrated. The potential break occurred between +260 and +-300 mv. vs. the saturated calomel electrode. After the uranium (IV) was oxidized, iron(1l) was titrated with dichromate. The potential break occurred between $660 and +700 mv. vs. the saturated calomel electrode. The results were accurate to 2 ~ 0 . 3 5 % . Interferences included titanium, vanadium, arsenic, molybdenum, tin, and antimony.
P
separation is generally necessary for the determination of iron in the presence of uranium or uranium in the presence of iron. Iron may be precipitated as the sulfide after the uranium is complexed Kith tartrate (23). Alternatively, iron may be precipitated and uranium complexed with ammonium carbonate (11). I n either case, reprecipitation of the iron is necessary to obtain accurate results. Several procedures are available if the uranium, rather than the iron, is precipitated. I n this case previous reduction of uranium(V1) to uranium(1V) ie recommended (2, 11). The reduction of acidic solutions of uranium and iron in a Jones reductor yields a mixture of uranium(III), uranium(IV), and iron(I1). Ewing and Eldridge (5) obtained three potential breaks when this mixture was oxidized with potassium permanganate. Uranium a t the 100-mg. level was accurately determined in the presence of 5.6 mg. of iron. Rodden (11) observed, however, that a n increasing negative uranium error occurred with increasing concentrations of iron. I n the titration of 0.36 gram of uranium in the presence of 2.0 grams of iron, the uranium error was 2.0%. Other metallic reducing agents have been used to determine uranium, inRlOR
cluding silver ( I ) , lead (7), cadmium amalgam (S), and liquid lead and bismuth amalgams (14). Lead quantitatively reduces iron(II1) to iron(I1) and uranium(V1) to uranium(1V) ; however, the use of lead in reductors was limited for many years because sulfuric acid reacted Kith the metal to form lead sulfate, which decreased the efficiency of the reductor. Cooke, Hazel, and AIcXabb ( 4 ) solved this problem by adding hydrochloric acid to sample solutions. This paper describes the reduction of uranium(V1) and iron(II1) by a lead reductor, followed by the successive oxidation of uranium(1V) and iron(I1) by potassium dichromate. The end points are determined potentiometrically. EXPERIMENTAL
Apparatus and Reagents. The lead reductor consisted of a glass column 35 cm. long and 2 cm. in diameter, with a reservoir 11 cm. long and 3.5 cm. in diameter a t the top of the column. The flow through the reductor was controlled with a 4-mm. straight-bore stopcock, 10 cm. long. A small glass-wool plug was inserted in the bottom of the reductor t o prevent lead from entering the stopcock. Small portions of reagent grade, granulated lead (hfallinckrodt Chemical Go.) which passed through a 25-mesh screen were washed R-ith 10% HCl and added to the reductor, which contained enough 10% HCl to keep the lead completely immersed. After each addition of lead, the column was gently tamped with a glass rod. When the lead in the reductor reached 25 cm., approximately 3 cm. of 5-mm.diameter glass beads were added to the top of the reductor. When the reductor was not in use, the lead was covered with a 10% HCl solution containing O.lyoferric ion. The samples were passed through the reductor into an 800-ml. beaker. The reductor, electrodes, thermometer, gas bubbler tube, and buret entered the beaker through holes in a polyethylene cover. The solutions were stirred and heated with a Temco Stir-Plate, Model SP-1025B. Potential readings were taken with a Beckman p H meter, Model 76. The electrode pair was a Beckman saturated calomel electrode (No. 39170) and a
Beckman thimble-type platinum electrode (No. 39271). Periodically, the platinum electrode was cleaned with scouring powder. Argon was bubbled through the sample solution by a sintered-glass Bter stick. A conventional 50-ml. calibrated buret was modified so that a horizontal elongation between the buret tip and the stopcock separated the main body of the buret by several inches from the sample solution. The modification minimized volume changes of the titrant while the sample was being heated. A water-jacketed buret served equally well. Standard Uranium Solution. National Bureau of Standards U108 was ignited a t 900' C. for 1 hour. Approximately 11 grams of the oxide was accurately weighed and dissolved in a minimum quantity of concentrated "Os, 20 ml. of concentrated H2SOI was added, and the solution was evaporated to SO* fumes. After the sides of the beaker had been washed with distilled water, the sample was evaporated again to SO, fumes, 7 ml. of concentrated HCl was added, and the samde was diluted to 1 liter with 3N HC1: Standard Iron Solution. The standard iron solutions were prepared by dissolving approximately 2.5 grams of accurately weighed iron (Jarrell Ash, 99.9573 in 250 ml. of concentrated HCl and diluting t o 1 liter with distilled water. Standard Potassium Dichromate. National Bureau of Standards potassium dichromate (99.98% oxidizing strength) was dried a t 135' C. for 1 hour. After cooling, 4.9045 grams was dissolved and diluted to 2 liters with deionized water (or copper-free distilled water). PROCEDURE
Dissolve the sample, containing from 1 to 935 mg. of uranium and from 10 to 550 mg. of iron, in a mixture of concentrated HC1 and "01. After dissolution, add 15 ml. of concentrated H$Oc and evaporate to SO* fumes. Carefully rinse down the sides of the beaker with distilled water and repeat the evaporation. Add 120 ml. of distilled water and 45 ml. of concentrated HC1. While the sample is cooling to room temperature, wash the reductor with three 25-ml. portions of 3N HCl and five 50-ml. portions of 1:15 HC1. VOL 35, NO. 3, MARCH 1963
309
Table 1.
Determination of Uranium and Iron
u,mg.
Fe, mg.
Error,
Found
Added
Found
U
Fe
0.9351 80.65 94.46 116.6 120.4 201.6 212.6 935.1
0.9309 80.77 94.40 117.0 120.5 200.9 212.5 936.7
550.0 139.2 86.99 278.4 147.6 55.68 164.3 5.500
550.9 139.5 87.07 278.4 147.2 55.60 164.0 5.445
-0.45 +0.15 -0.06 10.34 +0.08 -0.35 -0.05 +0.17
$0.16 $0. 22 +o. 09 0.00 -0.27 -0,14 -0.28 -1.00
Table II.
Effect of Various Cations on Uranium and Iron Results
U Fe U Fe Cation added, mg. added, mg. added, mg. found, mg. found, mg. 94.87 55,7l 2.5Ti 96.24 55.82 126.2 95.23 126.2 95.23 95.23 116.7 116.7 116.7 116.7 201.6 233.7 233.7 233.7 233.7
a
70
Added
71.13 55.74 71.43 55.74 55.74 71.13 71.13 71.13 71.13 55.68 137.5 55.01 55.01 55.74
2.5V 2.5Cr 2.5Mn 2 . 5 Co 2.5X 2.5Cu 100 Cu
1.0Cu" 2.5As 2.5Zr 2.5Mo 2.5Sn 2.5Sb 2.5Bi
132.4 94.93 126.1 95.11 94.93 116 . .. i ~~. 116.4 103.5 118.9 201.5 241.0 238 0 237.0 233.7
71.18 25.80 (1.31 55.63 55.65 71 ox 70.95 70.74 71.06 55.68 137.3 A4 . - 67 . 55.01 55.74
Error,
U
Fe
+1.44 $4.91 -0.32 -0.08 -0.13 -0.32 -n 17 -0.26 -11.3 $1.89 -0.05 $3.12 +1 , - 84 _+1.41
+0.20 $0.07 +O.ll
-0.17 -0.20 -0.16 -0 07
-0.25 -0.55 -0.10 0.0 -0.15 -0
0.0
A2
0.00.0
Interference added t o titration vessel and not passed through reductor.
Allow each portion t o drain t o the top of the glass beads before the next addition, but do not expose the lead t o the air. Place the polyethylene cover containing the electrodes, the thermometer, and the sintered-glass filter stick on a n 800-ml. beaker. Position the beaker on the stir plate below the reductor. Calculate the volume of 0.05N K2Cr207 required to titrate the estimated quantity of uranium in the sample. Transfer 95% of this volume from the buret to the beaker. (If the approximate uranium concentration must be determined, titrate one representative sample without prior addition of dichromate.) Add the sample to the top of the reductor and regulate the flow t o approximately 3 drops per second. Stir the solution and regulate the gas flow so that a steady stream of argon passes through the solution without spattering. Wash the solution through the reductor with three 25-m1. portions of 3N HC1, followed by five 50-ml. portions of 1:15 HCI. Heat the solution to between 85' and 95' C. Titrate the uranium with potassium dichromate delivered in 0.10ml. increments. Follow the potential on the p H meter. The uranium(1V) t o uranium(V1) end point occurs between +260 and +300 mv. us. the reference electrode. Equilibrium near the end point is slow; therefore, 2 t o 3 minutes must elapse between additions of titrant. (If the equilibrium time increases, clean the platinum electrode with scouring powder and wash thoroughly with deionized water.) Calculate the end
310
ANALYTICAL CHEMISTRY
point using the second derivative method (16). Continue the titration to determine the iron content of the sample. The iron(I1) t o iron(II1) potential break occurs between +660 and +700 mv. Add the titrant in 0.10-ml. increments in the vicinity of the end point. The potential break is small, but definite, Again calculate the end point by the second derivative method and subtract the volume of titrant required for the uranium titration. When the reductor is left overnight or longer, add 0.1% ferric chloride in 10% HC1 to the reductor. Unless this small amount of ferric ion is added before storage, the subsequent analysis yields low uranium results (4). RESULTS A N D DISCUSSION
The concentration of uranium was varied from 1 to 935 mg. and the concentration of iron from 0.5 to 550 nig., which are the practical limits for 0.05N KzCrz07. However, a t concentrations of iron less than 7 mg., the iron break and the normal potential drop between the iron and uranium breaks were difficult to resolve. Otherwise, results agreed within *0.35% of the theoretical values. The results of eight typical titrations are shown in Table I. In 12 successive titrations of 120.4 mg. of uranium and 147.6 mg. of iron, the uranium standard deviation was 0.30 while the iron standard deviation was 0.24.
Essentially all possible interfering anions (nitrate, fluoride, sulfide, peroxides) are removed by H2S04fuming and were, therefore, not investigated. Interfering cations were those which are reduced by a lead ieductor and subsequently oxidized by K2CrzO;. For interference studies, cations were prepared in their most stable oxidation state and then added to the uranium and iron mixtures. Results are shown in Table 11. Small amounts of copperiII) catalyze the oxidation of uranium(1V) to uranium(T'1) ( I S ) . Experimentally, a negative uranium error of 11.3% \vas obtained when 1.0 mg. of copper was added directly to the titration vessel (instead of being passed through the reductor). However, when the copper(I1) was passed through the reductor with the sample, reduction to copper metal occurred iiear the top of the column and 100-mg. quantities of copper(I1) could be tolerated without seriousl:; affecting the uranium results. Patterson (10) reported that large amounts of copper(I1) will eventually decrease the efficiency of the reductor and cause low results. The osidation of uranium(1V) t o uraniuni(V1) with dichromate is relatively slow, but ferric iron serves as a catalyst for the reaction (8). lt7hen uranium as titrated experinientally in the absence of iron, equi!ibrium in the vicinity of the end point vvas slow and uranium results were consistently low. In the presence of iron, howex-er, the potential readings readily stabilized and quantitative uranium and iron results were obtained. This indicated that the iron(I1) was oxidized to iron(IT1) by dichromate and the iron(II1) oxidized uranium(1T') to uraniuni(VI), producing an equivalent amount of iron(I1). After the ovidation of uranium(1V) was complete, the dichromate oxidized iron(I1) to iron(II1). Furnian and Schoonover (6) observed a similar phenomenon while titrating uranium with ceric sulfate contaminated with iron. TT'hen dichromate was not added to the titration vessel prior to heating the sample, the uranium results had a negative bias, even though the iron results were quantitative. This negative error was believed to be caused by air oxidation of uranium(1V) prior t o titration and was eliminated by adding 95y0 of the oxidant t o the titration vessel before the sample TWS passed through the reduction column ( 3 ) . Attempts to determine uranium and iron simultaneously with R Jones reductor also yielded consistently low uranium results. This was evidently caused by induced oxidation of uranium(1V) to uraniuin(V1) during the oxidation of uranium(II1) in the presence of iron. Oxidation of uranium(II1)
to uranium(1V) without formation of uranium(V1) was not achieved; air oxidation, direct potentiometric titration with dichromate, and addition of chromium(II1) followed by potentiometric titration of the resulting chromium( 11) all proved unsuccessful. Khen determinicg uranium in the presence of iron, therefore, the formation of uranium(II1) should be avoided by reduction with the less active metal, lead. The lead reductor is easily prepared, amalgamation is unnecessary ( 6 ) , and the complications caused by the evolution of hydrogen are not encountered (9).
LITERATURE CITED
(1) Birnbaum, Nathan, Edmonds, S.
M.,
IND.ENQ. CHEY.. ANAL.ED. 12. 155
(1940). (2) Blalock, T. J., U.S. Bur. Mines, Rept. Invest. 5687 (1960).
(3) Bricker, C. E., Sweetser, P. B., ANAL. CHEM.25, 764 (1953). (4) Cooke, W. D., Hazel, Fred, McNabb, W. M., Ibid., 22, 654 (1950). (5) Ewing, D. T., Eldridge, E. J., J. Am. Chem. SOC.44,1484 (1922). (6) Furman, N. H., Schoonover. I. C., I h d . , 53, 2561 (1931). (7) Koblic, O., Chem. Listy 19, l(1925). (8) Kolthoff, I. M., Sandell, E. B., “Textbook of Quantitative Inorganic Analysis,” p. 580, Macmillm, New York, 1952. (9) Xilkantan, P., Jayaraman, N., IND. EYG.CHEY.,ANAL.ED. 11, 339 (1939).
(10) Patterson, J. H., U. S. At. Energy Comm., ANL 5410 (1955). (11) Rodden, C. J., ed.-in-chief, “Analytical Chemistrv of the Manhattan Proi_.. ect,” Nationh Xuclear Energy SeriG, Division VIII, Vol. I, Chap. 1, McGraw-Hill, Xew York, 1950. (12) Scott. W. W.. “Standard Methods . of Chemical Andysis,” p. 467, Van Nostrand. New York. 1939. (13) Sill, C: W., Peterson, H. E., ANAL. CHEM.24, 1175 (1952). (14) Someya, Kinichi, 2. anorg. allgem. Chem. 152, 368 (1926). (15) Willard, H. H., Merritt, L. L., Jr., ~
Dean, J. .Ail “Instrumental Methods of Analysis, p. 419, Van Nostrand, Xew York, 1958. RECEIVED for review September 27, 1962. Accepted December 26, 1962.
A N e w Spectrophotometric Reaction of Beryllium M. CABRERA
A.
and T. S. WEST
Chemistry Department, The University of Birmingham, Birmingham 7 5, England
b A highly sensitive method i s described for the spectrophotometric determination o f beryllium in the range 0.01 to 7 p.p.m. The color reaction between the reagent Fast Sulphon Black F and beryllium i s specific a t pH 1 1 to 11.2 in the presence of EDTA and KCN among some 25 cations examined. Fluoride, citrate, tartrate, etc., d o not interfere and a 10-fold amount of phosphate can b e tolerated. The only cations which tend to interfere are chromium(ll1) and tin(ll), and these can be rendered harmless b y previous oxidation to their higher valency forms. The stability, sensitivity, and freedom from interference of the proposed method compare favorably with those of standard methods now in use.
I
N THE course of investigating the
complex formation between copper (11) ions and the bishydroxybisazo dyestuff Fast Sulphon Black F ((3.1. KO.26990) (2)
N
we have found that not only is the reaction very sensitive for the spectrophotometric determination of copper (11), but one of the very few cationic interferences is a positive one due to beryllium (1). Since this reaction was consistently observed and could not be
removed by the addition of fluoride to bleach the beryllium complex, we decided to investigate it further, in view of the limited number of satisfactory spectrophotometric methods available for beryllium. EXPERIMENTAL
Reagents. Fast Sulphon Black F, 1.55 grams of reagent in 1 liter of water. This solution diluted 10-fold provides the 2 x 10-4Llfreagent solution. Beryllium sulfate, 10-4M, prepared from a standardized 1Ow2Msolution of BeS04.4Hz0 containing 1.772 grams of solid per liter of slightly acidified water. Potassium cyanide, 10-zX, 0.65 gram of KCN per liter of HzO. EDTA (disodium salt), lO-ZW, 3.72 grams of disodium (ethylenedinitrilo) tetraacetate dihydrate per liter of HzO. Concentrated ammonia. sDecific zravity 0.88. pH 11 to 11.2 Buffer. Mix equal Darts of 10-lX NaOH and a solution which is 10-lM with respect t o both glycine and sodium chloride. Apparatus. Unicam SP.600 spectrophotometer with 4-cm. cuvettes. Procedure. CALIBRATION.I n t o 100-ml. standard flasks pipet 0 to 8 ml. of 10-4M beryllium sulfate solution plus 3 ml. of 10-*M K C N , 2 ml. of 10-2M E D T A , and 10 ml. of Fast Sulphon Black F reagent, followed by either buffer solution (5 to 10 ml. a s required) or 0.5 t o 1 ml. of NH3-i.e., sufficient t o obtain p H 11 to 11.2 as measured by a p H meter. Allow the color to develop for a minimum of 1.5 hours (preferably 3 hours if the p H 11to 11.2 buffer isused), dilute to volume, then measure the absorbance of the solution containing no v
beryllium against the test solution containing Be in 4-em. cuvettes a t 630 mP * The amount of cyanide may be increased as required and EDTA may also be added in amounts necessary to pvercome interference from foreign ions. With larger amounts of these masking agents, due heed must be paid to their influence on the p H of the solution. 1 ml. of lO-4-U Be + 2 solution = 0.902 pg. of Be+2 Determination. Proceed as above using near-neutral test solution, b u t add KCN and E D T A in larger amounts if necessary, with similar modification of the calibration procedure. NATURE OF COLOR REACTION
The color developed is similar to that with copper(II), suggesting, as would be expected for a small cation such as Be+2,that the color formation is due to configurational effects on the organic molecule rather than on the cation. The absorption curve of the reagent is shown in Figure 1, along with that of the same amount of reagent in the presence of 0.5 and 1 mole ratios of Be+2. It will be seen that the absorption band due to the beryllium complex a t pH 11 is of somewhat low intensity relative to the reagent band a t 600 or 450 mp. The complex is a 1 to 1 rather than 1 to 2 complex of beryllium with the reagent; further addition of beryllium produced only a slight change, presumably due to a mms-action effect. Subsequently, the application of a Job’s method of continuous variations revealed t h a t this was true and a 1 to 1 complex is formed. VOL. 35, NO. 3, MARCH 1963
31 1
