ANALYTICAL CHEMISTRY
1150 Table VII. Determination of Molybdenum in Plant Material by Thiocyanate-Stannous Chloride Method without Ether Extraction
Sample
Molybdenum In Original Sample, P.P.M.
Molybdenum Added, P.P.M.
Molybdenum Found, P.P.M.
Molybdenum Recovered, P.P.M. Yo
tained negligible amounts of molybdenum even with material high in silica; the molybdenum thiocyanate complex is determined directly in the hydrochloric acid solution without extraction; and nitrate and ferric ion are added to the unknown solution and to the standards. The method was tested by the addition of known amounts of molybdenum to plant samples before ashing. Table VI1 shows the molybdenum content obtained by the proposed method of the original samples as well as of samples to which molybdenum was added. The amounts recovered were nearly equal to those ad-led. LITERATURE CITED
Barshad, I., Soil Sci., 66, 187 (1948). Grimaldi, F. S., and Wells, R. C., IND. ENG.CHEW,ANAL. ED., 15,315 (1943).
tirularly in plant samples high in molybdenum where the possibility of incomplete extraction exists. The absence of the interfering ions, chromium, vanadium, tungsten, and rhenium ( 4 , IO), in solutions of plant ash also justifies the elimination of extraction. The concentration of molybdenum in solutions of plant ash encountered ranges approximately between 0.02 and 2.0 p.p.m. The sensitivity in the range between 0.02 and 0.20 p.p.m. is approximately 0.001 p.p.m. per 0.5% absorption and between 0.1 and 2.0 p.p.m. is approximately 0.01 p.p.m. per 0.5% absorption. The foregoing method differs from that proposed by RIarmoy ( 7 ) in that the filter paper and residue after filtration are not ignited and resulting ash fused with sodium carbonate, because spectrographic analysis and fusion revealed that the residue con-
Hiskev, C. F.. and Meloche. V. W., J . Am. Chem. Soc., 62, 1565,
isis (1940); 63,964 (1941).
Hoffman, J. I., and Lundell, G. E. F., J . Research Natl. Bur. Standards, 23,497 (1939).
Hurd, L. C., and Allen, H. O., IXD.ENG.CHEM.,A a . 4 ~ ED., . 7, 396 (1935).
Latimer, M .W., “Oxidation Potentials,” Sew York, PrenticeHall, New York, 1938. Marmoy, F. B., J . SOC.Chem. Ind.,58, 275 (1939). Rider and Mellon, IND.ENG.CHEM.,ANAL.ED., 18, 96 (1946). Sandell, E. B., “Colorimetric Determination of Traces of Metals,” New York, Interscience Publishers, 1944. Sandell, E. B., IND.ENG.CHEM.,ANAL.ED.,8, 336 (1936). Scott, W. W., “Standard Methods of Chemical A4nalysis,”New York, D. Van Nostrand Co., 1939. RECEIVED .4pril 26, 1948.
Use of Tributyl Phosphate for Separating Acetic Acid from Hydrochloric Acid H. ARMIN PAGEL, PAUL E. TOREN,
AND FRED
W. MCLAFFERTY
University of Nebraska, Lincoln, Nebr.
N .W earlier paper it was shown that the fraction of acetic Iphosphrtte acid extracted from aqueous solution by means of n-tributyl is essentially independent of the acid concentration. (1)
Hydrochloric acid, however, showed a pronounced increase of extraction with increased concentration; hence a reasonably complete separation of acetic acid from hydrochloric by this method appeared questionable if the concentration of the latter were fairly high. Work recently completed shows, hoxever, that the fraction of the hydrochloric acid extracted by the ester phase can be expressed by the equation:
in a constant temperature bath (25.00’ * 0.05” C.) and shaken vigorously for about 1 minute a t 15-minute intervals for several hours, without being removed from the bath. Complete phase separation took place after about 10 hours a t constant temperature, after which pipetted portions of each phase were analyzed. The hydrochloric acid was removed from the ester sample by extracting three times with about 10 ml. of water each time, followed with a fourth portion containing a small amount of base. The hydrochloric acid was then determined gravimetrically as silver chloride. The acid in the aqueous samples was determined by titration with carbonate-free sodium hydroxide. Where mixtures of hydrochloric acid and sodium chloride were
Table I.
This relation holds for mixtures of hydrochloric acid and sodium chloride as well as for hydrochloric acid alone, as shown in Tables I and 11. Therefore, if the hydrochloric acid in a mixture of the ti?-oacids is partially neutralized to a pH approximately the same as that of a comparable concentration of acetic acid alone, a satisfactory separation should take place. This wa- found to he true. PROCEDURE
In the work reported earlier, the amount of hydrochloric acid extracted into the ester phase was determined by draining the aqueous layer (lower) from a separatory funnel and then determining the acid in the ester by titrating with standard base. Because the hydrochloric acid concentration in the aqueous phase is always very high compared to that in the caster, significant errors resulted from the aqueous film left in the separatory funnel. The procedure was modified as follows: Large volumes (80 to 100 ml.) of ester and of aqueous solution were put into a 250-ml. glass-stoppered flask, which was placed
Distribution of Hydrochloric Acid in Two-Phase System Watern-Tributyl Phosphate
Hydrochloric Acid, .lf Ester phase Water phase
=
0,035 0.035 0.034 0.034 0.036 0.037 0.038
0,00390 0,00124
0.187 0 334 0.631 0,930 1.211 1.445 1.808
(CHCl)eater ( C H CCI-)wster
0.0134 0,0292 0.0518 0.0771 0,1257
Table 11. Distribution of Hydrochloric Acid in Two-Phase System Watern-Tributyl Phosphate, in Presence of Sodium Chloride Water Phase (Cl-), -21
(H+), M 0.0822 0.0970 0.442 0,446 0,874 0.879 0.884 0.886 0.890
1.078 1,105 0,688 1.460 1.877 1,126 1.918 1.398 1,017
Ester Phase (HCL), M ’
(CHCl)eiter = K ( c H + CCI-)water
0.00288 0,00348 0.00992 0,0203 0.0577 0.0331 0.0596 0.0418 0,0301
0.033 0,033 0.033 0.031 0,035 0.034 0.035 0.034 0.033
V O L U M E 2 1 , NO. 9, S E P T E M B E R 1 9 4 9 used, the totsl chloride was likewise determined gravimetrically. Sr'ns porous porcelain filtering crucibles were used. RESULTS
The data in Table I were used to calculate the distribution coeKcient. All equilibrium concentrations are calculated to mc:", per liter. Complete ionieatition of the hydrochloric acid in the water phase was assumed for calculating the constant in Tahle I. The ohloride concentrations in Table I1 are total chloride, and i t is assumed that both the acid and salt are completely ionized iii the water phase. The constants calculated for 'Ton products" ranging from about 0.035to 3.2 in Tahlcbles I and I1 arc in reasonably good agreement. In ordex to test the cornplcteness of separating acetic acid from hydrochloric acid by controlling the hydrogen and chloride ion products, the following experiments were performed: A mixture of aqueous acetic acid and hydrochloric acid was neutralized to pH = 1, using a Fisher titrimeter. The ester
22.
115s
was then added and the regular procedure was followed. Analyses of 20-ml. pipetted portions showed 0.351 M total chloride in the aqueous phase and 8.7 X 10-4 M hydrochloric acid in the ester phase. By calculation, the amount of hydrochlorio aoid in the ester phase is 1.12 X 10-8 if the value K = 3.4 X lo-* is used; hence the agreement is satisfactory. A seaond similar ,mixture was neutralized to pH 2 before extracting with the ester. Here the amount of hydrochloric acid from the ester sample gave only B faint qualitative test as silver chloride and was too small to he determined grwimctrically. Finally a mixture of 0.5 M acetie acid and 0.5 M sodium ohloride was extracted. In this ease a positive test for chloride from the ester sample was questionable.
It is logical to assume that other similar weak organic acids can be separated from hydrochloric acid by this method. LITERATURE CITED
(1) Pagel, H. A,. sod McLafferty, F. W., ANAL. C n e ~ . 20, , 272 ( 1948).
RECE~YE August D 4. 1948.
Uranyl Nitrate Hexahydrate, UO,(NO,),.GH,O
Excellent crystals of uranyl nitrate hexahydrate can be ohtained from either water or alcohol. Care must he taken to avoid crystallization of the anhydrous salt, becawe the hydrate is stable only at low temperatures. Either of these two phases may be ohtained on a microscope slide; the hydrate only in the presence of excem water.
Principal Lines Armour
d
I/Ic
d
Hanawslt
Armour
I/II
Hanawalt
Figure 1. Uranyl N i t r a t e Hexahydrate Left. Crystal%from water on a mioroaoope slide Right. Crystals from fusion, crossed Nicols
a
b
CRYSTAL MORPHOLOGY (determined by W. C. McCrone). Crystal System. Orthorhombic. Form and Habit. Plates and tablets usually elongated parallel to c. The crystals show hrachypinacoid ( 0101; macropinscoid ( 1 0 0 ) ; maorodome(011); and bipyramid ( 1 1 1 ) . Axial Ratio. a:b:c = 0.877:1:0.609; 0.874:1:0.609 ( 4 ) ; 0.8737:i:n.6nss
(e).
Intorfacid Angles (polar). 101 A io1 = 59O 30'; 011 A nil = 62" 40'. X-RAYDIFFRACTION DATA(determined by W. C. McCrone). Space Group. Tihl' (3). CellDimensiansb a = 11.58 A,, b = 13.20 A,, = 8.04 A. a = 1 3 . 1 5 A . , b ~ 8 . 0 2 A . , c = 1 1 . 4 2 A . ( S ) .a = 7 . 9 3 A . , b = 1 1 . 4 5 A . , e = 13.01 A. (1). Formula Weights per Cell. 4. Formula Weight. 502.18. Density. 2.73 (x-ray); 2.807 (1); 2.742 ( 3 ) .
Figure
2. Orthographic Projeotion of Typical Crystal of Uranyl N i t r a t e
Hexahydrate
OPTICAL PROPERTIES (determined by W. C., McCrone). Refractive Indexes (5893 A,; 25" C.). a = 1.482 + 0.002. B = 1.494 A 0.002. = 1.572 * n.002. 6 = 1.497(4).
