Engrnyring Process
development
Separation of Chromium Vanadium Salts by Liquidliquid Extraction ALBERT E. WEINHARDTI
AND A. NORMAN HIXSON UNIVERSITY OF PENNSYLVANIA, PHILADELPHIA 4, PA.
T h e separation of chromium and vanadium by the usual methods is difficult and, in commercial processes, expensive. This work was undertaken to develop a selective solvent, so that the separation could be accomplished by liquid-liquid extraction. Methyl isobutyl ketone has been found to be remarkable in preferentially extracting sodium dichromate from an aqueous hydrochloric acid solution of dichromate and vanadic acid. At low hydrochloric acid cmncen trations and in the temperature range from 0' to 25" C., separation factors well over 4000 were obtained. This selectivity indicates that methyl isobutyl ketone has considerable promise for industrial application, since solutions of chromium and vanadium compounds commonly result from the treatment of chromite or chromebearing ores. The very favorable separation indicates the possibility of using the extraction as an analytical method.
T"':
separation of chromium and vanadium has long been a problem of academic and comnlercial interest because of its very difficulty. The reported methods ( 6 , 11, 19, 29) are restricted t o use in analytical separations by the prohibitive cost of the reagents. They effect the separation by the precipitation of complex organic compounds of t,he metal ( 1 1 ) or by extraction of nietallic complexes or unstable compounds ( 6 , 19, 29). The specially controlled conditions lead to operating difficulties in any commercial use. A simple method of separat'ion such as liquid-liquid extraction would be economically attractive, especially if no compIexiiig agents were required. Although the separation of inorganic solutes by liquid-liquid extraction has received little attention, as far back as 1877 Skey ( 3 1 ) reported ethyl ether for extraction of certain metallic chlorides and sulfocyanates. Extraction of ferric chloride from aqueous hydrochloric acid solutions by ethers has received most attention. This system was first studied by Rothe ( 2 ? )in 1892 and had been studied recently b y many investigators ( 1 , 2, 12-14, 26, 23, 32, 3.4, 36). Of the many systems (S, 9, 16, 20, 22, 24, 26, 33)investigated, in most cases the inorganic solute was extracted from an aqueous hydrochloric acid solution by an ether. Hydrochloric acid seems singulir in its ability to increase the degree of extraction of inorganic salts from aqueous solutions. However, recent data ( 7 , 13') show that in ext,raction of metallic chlorides, the chloride ion concentration controls. In these cases, the hydrochloric acid was replaced in part or entirely by another chloride. Some investigators ( 7 , 9, 28) h a w considered the use of solvent's ot,her than ethers. Lacking applicable information in the literature, the problem of chromium and vanadium separation imniedint,ely resolved I
Present address, E. I. du Pont do Semours & Co., Kilmington, Del
itself into two rei:ited phases of prelirninary investigation: selection of the salts of chromium and vanadium to be separated, which obviously is dependent on the condition of the salt solution from which the extraction is to be made: and finding the sol.i.cnt that shows the greatest selectivity for one or the other salt. It was decided to limit the study to tlie highcr valence stxtos of chromium and vanadium-Le., CrC8 and I*+$. Those considered were the dichromatcv, or chromates, and the iuctavanadates. rlfter the system had becii conipletely defiried, the prol)lom became a detailed study of the equilibrium distribution 01' the salts between the organic and aqueous phases. MIATERIALS
Preliminary investigations were made with Baker's C.P. analyzed or Eimer and Aniend C.P. reagent grade salts. The sodium chromate, sodium dichromate, hydroc*hloric acid, arid sodium hydroxide contained no significant impurities in excess of 0.00570 and were used without purification. Anhydrous sodium chromate was prepared by heating overnight a t 150" C. Xellor ( 1 8 )reports no decomposition under 1000 C. S o analyses were presented for the sodium metavanadnte or ammonium metavanadate. h purification described by Liiigane and Meites ( 1 6 ) showed that the ammonium metavanadate as received was approximately 98% pure. The remainder \vas mainly ammonium chloride and a small amount of moisture. ?io further purification was necessary because the addition of hydrochloric acid to convert, the salt to vanadic acid formed ammonium chloride. The literature contained no method for the preparation of anhydrous sodium me1,avanadate. Successive determinations a t various temperatures indicated thst the salt could be dehydrated a t 120" C. The organic compounds used were of the best purity available from Fastman Kodak Co. and Paragon Testing Laboratories, and were used for the preliminary work as received. Thc methyl isobut,yI ketone used in the equilibrium studies \vas distilled in a 12-ball Snyder column and the heart cut h i l i n g between 115' and 116" C. was used. The refractive indcx \vm compared with the value 1.3955 a t 20" C. ( 1 0 ) and an)- batch showing a value outside the range of 1.3953 to 1.3957 at 20" C. \vas redistilled. PRELIMINARY INVESTIGATION
Method. Two methods of contacting the salt with thc solvr,iit were used. Initially some verification of the proposed rclal ionship ( 7 ) between anhydrous salt solubility and extractnbilit y as sought. To do this, a few grams of the anhydrous salt, v.-i:re added to 30 t o 50 ml. of solvent. Solubility was estiinated qualitatively from the amount of undissolved salt and the coloration of the solvent. No solubility was found with any of the solvents investigated. h second method was applied to a limited number of solvents. To a given amount of stock salt solution, enough hydrochloric acid or sodium hydroxide was added to make a solution 0.1 !If 1676
July 1951
1677
INDUSTRIAL AND ENGINEERING CHEMISTRY
TABLE: 11. QUALITATIVE EXTRACTIONS
TABLE I. ANHYDROUS SALTSOLUBILITY Solvent Alcohols Ethyl Isobutyl Capryl Cyclohexanol Diacetone Tetrafurfuryl Aldehydes Benzaldehyde Furfural Amines Methylaniline Pyridine Esters Amyl acetate Dibutyl phthalate Dibutyl tartrate Isobutyl lactate n-Butyl hydracrylate Methyl salicylate Methyl hydroxyisobutyrate Ethyl hydroxyisobutyrate Ethyl ethoxypropionate Monacetin Diacetin Triacetin Ethers Isopropyl n-Butyl Dichlorodiethyl Glycols Ethylene glycol Propylene glycol Ethyl hexanediol G 1y 6er o1 Diethylene glycol Triethylene gl col Dipropylene g~ycol 13-Butylene glycol duty1 Cellosolve Ethyl Carbitol Butyl Carbitol Glycerol monochlorohydrin Glyceryl monoricinoleate Polyethylene glycol, di-, triricinole:ate Diethylene glycol monolaurste Diglycol oleate Polyethylene glycol mono-oleate Hydrocarbons n-Heptane Benzene Styrene Ketones Methyl 'isobutyl Methyl n-amyl Methyl cyclopropyl Diisopropyl Acetophenone Benzophenone Cyclohexanone Aoetobutyrolactone Propiophenone Miscellaneous Ethylene chloride Dioxane Nitrobenzene n-Cresol Di-sec-amylphenoxy ethanol Octyl phenoxyethanol Octyl phendioxyethanol Octyl phentetraoxyethanol Triethyl phosphate
N a2C 04 Solubilitya
Reactionb
NS NS NS NS
NR NR NR NR R R
NS
9s
SR R
NS NS
SR NR
vss vss ss ss vss NS vss ss vss 9s vss vss
NR NR R R NR NR NR NR NR R NR NR
NS NS
NR NR NR
vs vs
vss vs vs vss vs
S S S
S S
S
S S S S S S S
R R
NR R
R
R R R R R R R R R R R R
NS
NR NR NR
NS NS NS NS
NR NR NR NR NR NR SR R NR
NS NS
vss NS vss vss NS
NS NS NS
5s
NS NS NS NS
vss
NR NR NR R NR NR NR NR NR
Nay08 Solubilitya NS SR Ethylene glycol NS NR Isopropyl ether NS NR Benzene NS SR Diisopropyl ketone NS NR Capryl alcohol NS SR Methvl csclourowl ketone . ~~. a VS, very soluble; S, soluble: SS, slightly soluble; VSS, very slightly soluble: NS, not soluble. b R, reaction; N R , no reaction: SR, slight reaction.
in salt and 1 M in acid or base. Exactly 100 ml. of either the acidic or basic salt solution were mixed with an equal volume of organic solvent and the distribution estimated. In sodium hydroxide solution the salts would be present as sodium chromate and sodium metavanadate, whereas in acid solution the chromates are converted to dichromates and the metavanadates to vanadic acid. The distribution was estimated from the coloration of the organic phase except for basic solutions of the vanadates which are essentially colorless. I n this case it was determined qualitatively by reaction with dimethylglyoxime. This method produced the solvent finally selected.
Results. The anhydrous salt solubility results are presented in Table I. Any reactions of the salt with the organic solvent are noted.
Solvent Ethers Is0 ropy1 n - g u t yl Dichlorodiethyl Alcohols Capryl Isobutyl Ketones Diisopropyl Methyl n-amyl Methyl ethyl Methyl isobutyl a
b
Extraction from Acid Solutionsa NazCrzOr HVOs Slight V. slight V. slight
V. slight None
Reaction Reaction Fair Good
*
Extraction from Basic Solutionsb NaaCrO4 KaVOa
Slight V. slight V. slight
None None None
Kone None None
Reaction V. slight
None None
None None
Reaction None None None
None None None None
None None None None
Acidic solution 1 M HC1. approximately 0.1 M solute. Basio solution,'l M N a O k ; 0.1 M NazCrzOr; saturated in NaVOs.
Table I1 summarizes the qualitative distribution data obtained by the second method. The solvents were selected on the basis of availability and cost. Discussion. Of the seventy-odd organic compounds of varied structures investigated only the polyhydric aliphatic alcohols showed any solubility for the anhydrous material. These compounds were ruled out by the oxidation-reduction reaction t h a t occurred and by their inherent water miscibility. Garwin and Hixson ( 7 ) found the anhydrous salt solubilities of cobalt and nickel chlorides to be a measure of their extractability by organic solvents. The discrepancy may be caused by the occurrence of chromium and vanadium as part of a complex anion, while cobalt and nickel were the cations. The second method was applied to ethers, alcohols, and ketones by selecting two or three representative compounds. Although diisopropyl ketone and methyl a-amyl ketone both gave negative results, a third ketone, methyl isobutyl, gave a distribution between aqueous acid solution and the organic solvent t h a t distinctly favored the organic phase. This distribution became more pronounced with increasing aqueous hydrochloric acid concentration and was selective for sodium dichromate, little or no vanadic acid being extracted. Substitution of sulfuric acid for the hydrochloric acid reduced the extraction tremendously. None of the solvents tried showed any appreciable extraction with sodium hydroxide solutions. Methyl isobutyl ketone is available in large quantities a t 12 cents per pound. Its water solubility is 2% by weight a t 20' C. EQUILIBRIUM DISTRIBUTION STUDIES One hundred milliliters each of diluted standard stock solution and of methyl isobutyl ketone were adjusted to temperature separately and then mixed in a 250-ml. flask. The flask was rotated 25 minutes in a bath. Five minutes were allowed t o ensure complete phase separation ; then duplicate samples of each phase were rapidly analyzed. The low limit of solubility of vanadic acid in dilute acid solutions prevented preparation of concentrated stock solutions A given amount of ammonium metavanadate was weighed into a 250-ml. volumetric flask and sufficient hydrochloric acid added to convert the salt to vanadic acid and to make the solution up to the desired hydrochloric acid molarity. The phases were contacted and sampled as before. Determinations were ma$ at 25" and 0" C. The temperat y e was maintained a t 25 * 0.1" C. by a water bath and a t 0 * 0.2 O C. by an ice bath. I n the system containing both sodium dichromate and vanadic acid, it was again necessary to weigh in a given amount of ammonium metavanadate. A measured volume of sodium dichromate stock solution was added with sufficient standard hydrochloric acid t o convert the ammonium metavanadate t o vanadic acid and to make the solution a definite molarity in hydrochloric acid. The resulting solution was made u to 250 total ml. with distilled water. A sample was analyzed i% Cr6 and V5 by titration with ferrous ammonium sulfate t o verify the over-all composition. One hundred milliliters were equilibrated with the ketone as before. After 30 minutes, four samples were withdrawn from each phase for analysis.
INDUSTRIAL AND ENGINEERING CHEMISTRY
1678
SODIUM DICHROMATE-HYDROCHLORIC ACIDWATER-,METHYL ISOBUTYL KETONE
Analytical Methods. AQUEOUS PHASE. A 10- t o 25-ml. sample was withdrawn from the aqueous phase a t equilibrium and immediately added t o a 260-ml. beaker containing 100 t o 125 ml. of distilled water. The water served the dual purpose of reducing the acid concentration and providing a convenient volume for analysis. The immediate reduction of the acid concentration was important because above 3 N hydrochloric acid an oxidation-reduction reaction between aqueous dichromate and hydrochloric acid takes place a t room temperature. Sodium dichromate was titrated with standard ferrous ammonium sulfate solution. A Serfass electron ray amembly (90) was required to indicat,e the end point accurately because of the highly colored solutions. T o check the reported accuracy of 0.1% in a titration requiring 30 to 40 ml. (4, 5 ) , a standard solution of potassium dichromate was analyzed. The results presented below show the accuracy t o be as good as reported.
60 ml.
Calcd. Amount of KzCr207 Present, Meq. 5.000
25 rnl.
2,500
60 ml. plus 5 ml. concd. HC1
5.000
5.002 5.003
50 m1. plus 15 ml. concd. HC1
5.000
5.000 4.997
50 ml. plus 1 5 ml. concd. HC1"
5.000
Aliquot Portion of Standard Solution Used
Amount Found by Titration, Meq.
4.996 5.004 4.998 2.504 2.501 2.499
5.004
4.9fl5 a Analyzed after 1 hour to determine if appreciable chlorine was formed. The presence of any significant amount of chlorine would have given erratic results.
The electrode systems a plicable to dichromate-ferrous salt titrations are the polarise{ platinum-platinum electrodes or the self-polarizing tungsten-platinum syst,em. The systems are equivalent. KETONEPHASE,The dichromate was extracted from the ketone phase by three separate extractions with distilled water or until the ketone gave no test with diphenylcarbohydrazide. The extracts were pooled and analyzed as described above.
Results. Equilibrium distribution determinations were made at two temperatures over a range of sodium dichromate and hydrochloric acid concentrations. The data are grouped according t o the original aqueous dichromate concentration-that is,
Figure 1. Effect of Hydrochloric Acid Concentration on Distribution Coefficient
Vol. 43, No. 7
prior t o addition of the organic solvent. The original dichromate concentration was maintained eonstant throughout a series Table I11 shows the effect of varying the initial hydrochlor~c acid concentration while maintaining constancy of initial salt concentration. Figures 1 and 2 show a plot of the log of the distribution coefficient, K , against the acid concentration. Figure 3 shows the decrease of the distribution coefficient with increasing dichromate concentration a t constant hydrochloric acid concentration. The relationship a t both 25" and 0" C takes a hyperbolic form indicating a possible limiting solubility in the ketone. Figure 4 (25' C.) is a logarithmic plot of the dichromate concentration in the ketone phase against the aqueous dichromate concentration a t equilibrium. Curves of constant original dichromate concentration and constant original acidity are shown. It is possible to estimate from this figure how a solution with a given original dichromate and acid concentration will distribute itself between the phases. Figures 5 and 6 present logarithmic plots of K against the Cre concentration in the two phases. They emphasize an interesting trend in the dilute regions and permit the estimation of K , the distribution coefficient, from a knowledge of the Crb concentration in either phase and the original aqueous hydrochloric acid concentration. Table IV shows the effect on the distribution coefficient of varying the volume of the organic solvent used in the extraction By its definition, the distribution coefficient will be affected by a variation in the volumes contacted. The true thermodynamic distribution coefficient is the ratio of the activity of the salt in the organic phase t o that in the aqueous phase. In the absence of activity coefficient data, this is often simplified to the ratio of the mole fractions. It can be reduced to the ratio of the molar concentrations in extremely dilute solutions, but only if several assumptions are made. Since the salt concentrations used were not in the range normally considered dilute, thc variation of the distribution coefficient with the ratio of the volumes contacted is not surprising. Discussion. In the range of dichromate concentrations investigated (0.02 t o 3.0 M ) the distribution coefficient is increased tremendously by increasing the aqueous acid concentration from 1 to 5 M . No distribution of sodium dichromate was observed
Figure 2. Effect of Hydrochloric Acid Concentration on Distribution Coefficient
Figure 3. Effect of Sodium Dichromate Concentration on Distribution Coefficient
INDUSTRIAL AND ENGINEERING CHEMISTRY
July 1951
1629
OF SODIUM DICHROMATE BETWEEN METHYL ISOBUTYL KETONE AND WATERIN TABLE111. DISTRIBUTION OF HYDROCHLORIC ACID
25' C -. Original Aqueous Molarity HC1
,
E uilibrium Concn., %illimoles/Liter Aqueous Ketone phase phase
ICa
00
c.
E uilibrium Conon., %illirnoles/Liter Aqueous Ketone' phase phase
20 MILLIMOLES O F NazCriOr 10.83 1.91 1.388 5.67 ......... 1.378 ...... ......... 1.357 ......... 1,402 1 5 . 5 1 6.22 0,504 2.494 15.49 6.19 0.511 2.502 14.42 6.25 0.497 2.307b 14.99 6.45 0.506 2.3291 ......... 0.505 ...... 16.09 11.25 0.3110 1.431 ......... 0.3128 ...... 16.92 16.78 0.2206 1.008 17.04 17.14 0,2194 0.994 14.99 22.94 0.1512 0.654b 17.00 22.66 0.1489 0.750 16.84 30.54 0.1149 0.551 16.96 30.50 0.1126 0.556 12.22 3 1 . 4 0 0.3891 C ..... 1 6 . 4 0 39.33 ..... 0.4170 15.82 3 9 . 5 5 . . . . 0.4000b 15.82 3 9 . 5 5 ..... 0.4000b 0.3184 16.16 5 0 . 8 . . . . 16.05 5 0 . 3 ..... 0.3191 0.3312 16.44 4 9 . 6 ..... 13.76 5 0 . 0 ..... 0,2751C 16.01 62.2 . . . . 0.2751 16.18 61.5 ..... 0,2632 14.57 7 2 . 9 . . . 0,1998 14.24 7 2 . 5 ..... 0.1962
Ka
I.
0.5
......
...... 1.0
1.5 2.0 2.5 3.0 3.5 4.0
4.5 5.0
0.5
OF NaiCrz07 11. 100 XILLIMOLES , . ,, ..,,. 39.54 ......
...... ...... ......
1.0 1.5 2.0 2.5 3.0 3.5 4.0 4.5 5.0
22.98 23.12 21.236 12.18 12.03 7.98 7.80b 5.47 5.48 4.184
...... 3,074 3.068 2.361 2.364 1.820. 1.851 1.7726 1.514
. . . . . . . . . . . . . . . . . .
. . . . . . . . .
69.8 71.4 65.0 81.4 81.4 85.1 82.0 87.1 87.2 89.4
3.04 3.09 3.06 6.68 6.77 10.66 10.51 15.92 15.89 21.36
89.3 89.2 89.6 89.3 87.0 87.7 84.2 85.7
29.06 29.09 37.95 37.78 47.8 47.4 47.5 56.6
.........
.........
18.45 18.47 18.45 18.54 19.30 19.42 19.34 19.43 19.43 19.46 19.52 19.72 19.74 19.73 19.69 19.83 19.80
.....
94.3 94.4
53.1 53.0
....
.....
94.3 94.5 96.6 96.8 97.1 97.2
105.3 104.5 147.0 147.2 186,Z 187.3
..... ..... ..... 0.3535 0.3518
3.0
......
4.0
5.0
55.4 57.3 57.3 55.3 31 44 32 90 32 58
.... ....
....
.... .... 98.6 98.1
..... .....
.....
1.0
4.0
92 1 895 710 540 540 369 6
4.5 5.0
312 8 256 2
2.0 3.0 *
Eq$,libpium Conon., illimoles/Liter , A ueouS Ketone &as, phase
500 MILLIMOLES OB 128.9 0.349 131.1 0.333 129.5 0.350 268.3 1.311 270.3 1.330 349.6 3.59 351.8 3.63
158.6 262.9 280.1 288.2 287.8 292.2 291.6 285.5
0.991 8.31 22.32 44.7 44.1 79.6 81.4 80.1
168.8 316.7
i?i:4
'0:5i7 0.516 2.16
48.7 49.2 47.5 25.10 24.52 24.55
390.1 389.8 385.7 412.4 411.7 402.3
8.01 7.93 8.10 16.43 16.79 16.40
13.08 13.38
418.9 419.1
32.03 31.32
..,.,
333.4 327.3 146.9
.....
.........
.....
.........
243.1 190.6 191.3
.
. . . . . . . . .
1000 MILLIMOLES O F NaiCrzOi 158.2 0.172 846 158.4 0.177 846 317.6 0.447 660 454 0.840 486 456 0.845 476 582 1.49 309.9 . . . . . . . . . 301.0
1.86 2.41
........
.........
.....
6.89 6.77 6.78 6.83 12.30 11.98 12.01
583 616
KQ
NazCrzOr
.........
381.9 387.8 388.5 377.7 386.9 394.3 391.2
V.
. . . . . . . . .
....
0.896 0.904 0.658 0.656 0.521 0.519
2.0
369.4 393.7 369.6 204.6 203.8 97.5 96.9
. . . . . . .
.....
.....
1.0
Ka
c-.
00
7 -
111. 300 ~ I I L L I M O LOF E SNazCrlOi 160.0 119.3 0,603 197.8 3.15 31.65 72.2 227.3 7.52 12.55 34.34 258.1 14.58 6.45 18.92 275.9 18.94 275.6 14.55 6.53 10.73 271.4 25.29 3.669 10.58 269.2 25.44 3.584 10.85 270.2 24.90 3.566 10.94 272.7 24.93 ..... IV.
. . . . . . . . . . . . . . . . . .
1.51 1.51 1.51 1.51 14.90 14.89 14.46
.....
5.0
. . . . . . . . . . . . . . . . . . . . . .... . . . . . . . . ... ... ... ...
59.8 58.7 58.8 59.7 90.4 90.2 89.6
1.776 1.782
1.0 2.0 3.0 4.0
13.29 13.40 13.60 13.22 38.26 37.99 38.90 38.37 38.49 62.6 62.4 89.4 89.9 130.5 132.2 172.6 175.8
38.81 38.99 39.48 6.06 6.06 6.20
.....
-
25" C-. E uilibrium Concn., %illimoles/Liter Aqueous Ketone phase phase
7 -
original Aqueous Molarity HC1
PRESENCE
THE
195.1 195.1 362.1 505 495 600 602 627 662 663
0.231 0.231 0.549 1.038 1.041 1.94 2.00 2.58 3.47 3.46
VI.
OF NazCr207 2960 MILLIMOLES 1.0 2870 275.7 0.0961 2890 320.4 0.111 2.0 2818 ,519 0.184 2820 568 0.201 2781 513 0.185 2807 563 0,201 3.0 2749 724 0.263 2755 766 0,278 2740 714 0.260 ..... ......... 4.0 2673 905 0.339 2675 947 0.354 a K ,millimoles per liter, organic phase/millimoles per liter, aqueous phase. b Contacted for 1 hour. C Contacted for 2 hours.
..... ..... 278.9 278.9
in the absence of hydrochloric acid, b u t the presence of even a small amount of the acid causes a distribution. Apparently the distribution coefficient varies continuously from zero a t zero hydrochloric acid concentration. Figures 1 and 2 also show t h a t the distribution coefficient increases most rapidly when the acid is increased from 0 to 1 M , especially in the dilute dichromate ranges. The increase of the distribution coefficient with a decrease in temperature from 25" to 0' C. is striking. It is greatest a t the lower dichromate concentrations where it amounts to 500 toBOO%, benewhile even at the highest concentrations investigated, ficial effect of some 3% was observed. The range of acid concentrations investigated was limited by the reaction of sodium dichromate with hydrochloric acid a t high concentrations. At acid concentrations below 3 M, the reaction in extremely slow at 25" C. and virtually nil at 0 " C. This appears t o be the best acid concentration range regardless of
temperature. Although the reaction increases in rate with acid concentration, there was no serious difficulty below 5 M at either temperature. It was necessary, however, t o assume t h a t equilib-
TABLEIV.
EFFECTON
THE
DISTRIBUTION COEFFICIENT OF
VARYINU THE VOLUMES CONTACTED Initial Vol., M1. Aqueous Ketone phase phase
Ratio of
E uilibrium Concn., rSlillimoles/Liter Ketone
Volumes
phase
phase
K
Initial Conon. of NazCriOr,100 Millimoles; Original Aqueous Molarity HCl. 1 M; Temp., 25O C. 25 0.25 61 1 147 0 2.41 100 100 50 0.50 41.0 109 6 2.67 100 75 0.75 29.78 85.4 2.87 100 100 1.00 22 98 69 8 3 04 100 1.33 17.35 55 7 3.21 75 50 100 2 00 11.42 38 94 3 41 25 100 4.00 5.26 20 14 3 83
INDUSTRIAL AND ENGINEERING CHEMISTRY
1680
I
Vol. 43, No. 7
I
CONCENTRATION AQUEOUS PHASE, MILLIMOLS C.-*/LITER
Figure 4. Variation of Distribution with Hydrochloric iicid and Sodium Dichromate Concentrations Figure 5. Variation of Distribution Coefficient with -4queous Sodium Dichromate Concentration
riuni vr-as attained even though the reaction was still proceeding. A reasonable contact' time was adopted and samples of the phases analyzed and reported as equilibrium concent,rations. Several determinations were made at longrr contact times. These results are discussed below. The following equation is usually considered representative of the reaction: 14HCl
+ NazCr2O7 = 2KaC1 + 2CrCI3 + 3C1, + 7H:O
Although chFomic chloride is not extractable, the formation of any appreciable amounts would effectively reduce the dichroinak concentration and irdd another electrolyte to the aqueous solution. This would be a problem only a t high acid concentrations where the reduct,ion of the total dichromate should cause an increase in the distribution coefficient. The presence of chromic chloride, sodium chloride, and chlorine produces a n unknown effect. The data in Table 111 marked with superscripts 6 and c show that allowing the reaction t o proceed further by prolonged contact reduces the amount of C Y found in both phases; the distribution coefficient remains constant. The presence of chlorine in the aqueous solution indicates the possibility of chlorinating the ketone. This apparently takes place to a negligible extent,. Rice and Fryling ($6) reported a. velocity constant of 18 X l o p 4nioles per minute for the chlorination of methyl isobutyl ketone b y an aqucous acidic chloride solution, 0.1 M in chlorine a t 25 C. The temperature coefficient of this chlorination was large indicating that a reasonably low temperatsure could be used a t which no chlorination would take place. The following refractive indexes show that essentially no chlorination has taken place. 1J.ater sitturat,ion does not change the refractive index appreciably. Material Methyl isobutyl ketone Pure anhydrous Satuiated with water After use. saturated with water Literature value Methyl dichloro isobutyl ketone, literature value Methyl tricbloro isobutyl ketone, literature value
Refractive Index at 20' C. 1.3936 1,3962 1.3932 1 . 3 9 5 5 ( 10 ) 1 , 4 6 3 2 (3') 1.4872 (5)
In Figure 3 the distribution coefficients decrease with increasing dichromate concentration and tend to converge, regardless of acid concentration. Thiu apparent convergence takes place when the aqueous phase approaches saturation in dichromate. Extrapolation indicates t,hat a t saturation t'he variation of hydrochloric acid concentration might have little or no effect. This is apparently t.he case, since the slopes of Figures 1 and 2 (log of the distribution coefficient against the molar acid concentration) decrease with increasing dichromate concentration. Bnalogous effects of hydrochloric acid have been observed in many systems. However, in all cas= the distribution coefficient increased with hydrochloric acid to a maximum value after which
further increase in hydrochloric acid concentration produced a decrease in the coefficient. No such maxima were observctl here; however, it may be evident a t extremely high acid concentrations. Lingane and Meit,es (16) first observed the marked effect or temperature on the distribution coefficient. They found that the distribution coefficient for the extraction of vanadic acid from aqueous solution by isopropyl ether increased approxirnatcly 25% when t'he temperature mas decreased from 25" to 0" C. Nachtrieb and Conway ($2) ob100'- , -- - d TEMF? 2 5 ' C . served over t h e same temperature. range about a 30% increase in the distribution coefficient for ferric chloride extraction by isopropyl ether. The effect with sodium K dichromate is more pronounced. I: e w data are available on the offect of salt -concentrrtt ion on the dit: t r i b u t io LL coefficient in t h v eo 11 c e n t r a t e t l rcgions; h o we v B I' 0.1 Vilbrandt and , I , , , 100 10 1000 McCormack (34)) CONC, KETONE PHASE,MILLIMOLS C : ~ L I T E R working with a.1Figure 6. Variation of Distribum o s t s a t u r a t e tl tion Coefficient with Sodium solutions of ferric Dichromate concentration in chloride, found thnt Ketone the d i s t r i b u t i o n coefficient incrcased t,o a maximum and then decreased sharply with in( concentration, an effect similar to t,he one observed in this work. The concentrations investigat,ed here are higher than those normally assumed to be in the range of applicability of the simple distribution law. However, the trends a t the log-er concentrations are such t h a t several interesting conclusions may be drawn from a consideration of Figure 1. The slope of the curves is approximately unity a t the loir-er concentrations in conformance with the law. Also, it approaches unity over a greater conceutration range a t higher acid concentrations. Sincc satisfact,ion of simple distribution law, presumes that t,he solute has the saine molecular weight in both Polventa, it, appears that the dis~
1
1
,
,
,
I
,
INDUSTRIAL AND ENGINEERING CHEMISTRY
July 1951
tributed species do not vary as rapidly with dichromate concentrations at high acidity as at low. T o interpret the curvature of the plot, it is necessary t o assume a partial association or dissociation. If either takes place in either or both phases, the curvature could be the result of the variance of the degree of association or dissociation. At the highest dichromate concentrations the slope reverses upward toward unity, tending toward the same effect noted in the dilute regions.
1681
1. Low acid concentration (less than 3 M ) 2. Dilute dichromate concentrations (0.02 to 0.3 M ) 3. Low temperatures (approximately 0" C.) VANADIC ACIWHYDROCHLORIC ACIDWATER-METHYL ISOBUTYL KETONE
Analytical Methods. AQUEOUSPHASE. The sampling procedure was identical with that used in the sodium dichromate determinations. The vanadium was determined by reduction with standard ferrous ammonium sulfate solutions using the Serfaas electron ray titrimeter reported by McNabb (17). The accuracy of the determinations is reported to be between 0.1 and 0.2%. 1 TEMF! Ooc
- -
IO
I--
K I
=--
--
0.1
2 0 UM HVO
Y
0.01-loo/:
300 I
2
Figure 7. Effect of Hydrochloric Acid Concentration on Distribution Coefficient
Figures 5 and 6 emphasize the trends previously observed in Figure 4. A plot of the log of the distribution coefficient against the log of the concentration in either phase should show a horizontal line in the regions of dilute dichromate concentration, if the simple distribution law holds. While Figures 5 and 6 show no well defined horizontal line, thc re is evidence that with an extrapolation t o more dilute couccntrations such a line would be formed. Again in the more concentrated regions the distribution coefficient decreases and because the necessary data t o permit the calculation of the thermodynamic distribution coefficients are lacking, it can only be hypothesized t h a t this decrease is due t o a change of solute species. Table IV shows the variation of the distribution coefficient with the volumes of the solvents contacted. The distribution coefficient, as the ratio of the concentration in the two solvents, should be invariant only in the extremely dilute region. The value of the distribution coefficient increased as the ratio of ketone to aqueous phase increased. Aluminum chloride, calcium chloride, and sodium chloride were tried but were ineffective as distribution promoters, in the absence of hydrochloric acid. The substitution of sulfuric acid for the hydrochloric acid was investigated. The data below indicate t h a t it is also ineffective. 2 5 O C.: 20 Millimoles
of NanCrnOr 2 M Ha904 2 MHCl
K 0.104
16.96
T h e optimum conditions for the extraction of sodium dichromate from aqueous hydrochloric acid solutions by methyl bobutyl ketone are:
MOLARITY OF AQUEOUS
HCI
Figure 8. Effect of Hydrochloric Acid Concentration on Distribution Coefficient
This was further substantiated by preparing a standard solution of V5 by dissolving a weighed amount of ammonium metavanadate in 6 N sulfuric acid. It was analyzed by reduction with standard ferrous ammonium sulfate, and also by reducing the V6 to V4 with sulfur dioxide, boiling off the excess sulfur dioxide, and titrating with dilute potassium permanganate to a persistent light pink color. The results tabulated below show t h a t the methods give results which check within 0.1%. I n addition, data are included which show that large amounts of hydrochloric acid do not affect the accuracy of the method employing ferrous ammonium sulfate. Aliquot Portion of Standard Solution 50 ml.
Millimoles of V5 Found KMnO4 0.9961 0.9949 0,9946 50 ml. Ferrous ammonium sulfate 0.9944 0.9953 0.9945 Fcrrous ammonium sulfate 50 ml. plus 5 ml. concd. HC1 0.9947 0.9953 50 ml. plus 20 ml. eonod. HC1 Ferrous ammonium sulfate 0 9952 0 9941 0.9944 50 ml. plus 20 ml. concd. HC1" Ferrous ammonium sulfate 0 9941 a Analyzed after 1hour to determine if any appreciable amount of chlorine waa formed. The presence of appreciable chlorine would have caused erratic results. Method
KETONEPHASE. The sample was withdrawn and extracted three times with an equal volume of water. The washings were pooled for analysis as with the aqueous phase.
Results. A series of equilibrium distribution measurements was made at two temperatures over a range of vanadic and hydro-
1682
OF VANADIC ACIDBETWEEN METHYL TABLE V. DISTRIBUTION ISOBUTYL KETONEA N D WATERIN THE PRESENCE OF
HYDROCHLORIC ACID
250 c -. Equilibrium Concn., Millimoles/Liter Aqueous Ketone phase phase
I. 1.0 2.0 3 0
4.0
5,0 6.0
20.04Q 19.550 19.01 19.09 19.09 19.10 17.46 17.42 17 O l b 12.85 12.86 12.42b 5,32 5,35 11.
2.0 3.0 4.0 5.0
6.0
7.0
2.0 3.0 4.0 5.0 6.0
96.8 96.8 95.2 95.1 94.6a 89.6 89.1 69.5 68.6 69.5 67.9b 32.54 32.47 8.44 8.50
00 c.Equilibrium Concn., Millimoles/Liter Aqueous Ketone phase phase
r
7 -
Original Aqueous Molarity HC1
K
20 hfILLI?IiOLES O F HJ'Os .... 0.0529 0,00264 0,1584 0.00810 .... 0 , 4 9 8 0,0262 19.40 0.525 0,0275 19.32 0 . 5 1 2 0.0268 .... 0.551 0.0288 .., . 1.791 0,1030 17.87 1.743 0,1000 17.82 1.727 0 1015 .... 5.75 0.448 13,62 5.79 0.450 13.58 5,60 0.451 .... 12.55 2.36 5,46 12.64 2.36 5.55 100 MILLIMOLES OF IIVOa 0 . 3 7 5 0,00387 .... 0,364 0.00376 .... 1 . 5 5 1 0,0163 95.2 1.506 0.0158 95.0 , . . 1.523 0.0161 0.0678 90.7 6 076 90 6 6 . 0 6 8 0.0681 0.327 22.75 22.96 0,335 7219 22.80 0.328 73.0 22.54 0.332 .... 55.4 1.70 35.93 55.1 1.70 36.16 75,8 8.98 9 78 75.5 8.88 9.74
111. 300 ~ ~ I L L I I I O L E O S F HVOs 292.0 0.911 0.00312 .... 285.7 3 . 6 6 1 0.0128 288.2 285.9 3 . 7 8 0 0.0132 288.2 267.2 16.70 0,0625 275.2 268.6 16.50 0.0814 275.0 264.lC 16.35 0,0619 .... 210.0 60.9 0.290 226.9 212.2 60 4 0.284 226.8 212.9 62.2 0.292 .... 1.16 139.1 121.0 139.8 139.2 139.6 1 . 1 5 120.9
0.463 0.462
...
7
K
... 1.716 0.0960 1.700 0 0954 6.19 6.14
0.464 0.452
. . . . . . . . . 14.48 14 74
2 65 2 66
1.390 0.0146 1.381 0.0145
. . . . . . . . . 5.540 5.664
0.0611 0.0625
....
23:58 23.45
0.323 0.323
63.16 62.95 92.4 92.3
1.76 1.74 9.44 9.47
.........
3.14 3.24 14.98 15.07
0.0109 0.0113 0.0544 0.0548
60.8 60.7
200
480
Concn. Aqueous Phase 140.5 350.0
0.268 0.268
162.6 152.5
1.10 1.10
Concn. Ketone Phase 43.40 95.2
K 0.309 0,272
Vanadic acid concentrations are in millimoles per liter. These data are plotted on Figure 7 with the data of Table V. Q
1
,
I l l
I
L
100
10
Figure 9. Variation of Distribution w i t h Hydrochloric Acid and Vanadic Acid Concentrations
. . . . . . . . .
COEFFICIENT AT VANADIC ACID TABLE VI. DISTRIBUTION CONCEKTRATIOKS O F 200 A N D 480 >IILLIMOLES PER LITER Origins! HC1 Acid Concn. 5 M 5 M
OIL
CONC. AQUEOUS PHASE, MILLIHOLS ;;/,ITER
chloric acid concentrations limited only by the solubility of vanadic acid in aqueous hydrochloric acid solutions. The data of Table V show the effect of varying the hydrochloric acid concentration a t 25" C. and 0 " C. for three initial aqueous vanadic acid concentrations. Figures 7 and 8 represent these data as log of the distribution coefficient against original hydrochloric acid concentration. Figure 9 shows the log of the vanadic acid concentration in the ketone phase against the log of the concentration in the aqueous phase. Table V I presents the data of two determinations a t 5 M
Original Vanadic y d Concn.
hydrochloric acid concentration and 25 C., but a t vanadic acid concentrations other than those investigated in detail. From Figures 7 and 9, it is evident that the distribution coefficient increases rapidly with increasing hydrochloric acid Concentration, although the values are small. Only when the aqueous acid concentration is increased to 6 M does the distribution coefficient exceed one. The acid concentrations of interest for chromium separation are 0 to 5 M . The higher acid concentrations were investigated merely t o determine the trend of the curves.
0.0238 0.0239
These results were obtained bv uoolina .. - several samules and the accuracy is open t o question. b Analyzed after 2 hours of contacting. 0 Analyzed after 1.5 hours of contacting.
Temp., C. 25 25
Vol. 43, No. 7
INDUSTRIAL AND ENGINEERING CHEMISTRY
In contrast t o the effect observed with sodium dichromate, temperature appears to have a reverse effect in thc acid range 0 to 5 iM. Above 6 M hydrochloric acid, however, a decrease in temperature produces an increase in the distribution coefficient. The reverse effect observed a t the lower acidities could be caused by a decrease in the mutual solubility of the phases. Figure 9 shows essentially straight lines of slope approximately one, although with dilute concentrations an upward curvature is apparent. This appears to be contradictory to the simple distribution law which would indicate that a slope equal to one should be observed in the dilute solution region. The absence of activity data and information on the solute species in the two solvents make interpretation of this trend impossible. Vanadic acid reacts with hydrochloric acid in a manner analogous to sodium dichromate. 2HV03
+ 6HC1
= 2VOC12
+ Cl? + 4H20
Lingane and &kites (16)found that this reaction proceeded slowly in 8 M aqueous hydrochloric acid. They observed that the per cent of the original amount present in the aqueous solution extracted by isopropyl ether decreased with time because of the reduction of the vanadate by the hydrochloric acid. They also found that there was no appreciable extraction of the vanadic acid until the aqueous hydrochloric acid concentration exceeded 5 M. They suggested that the solute extracted was probably an uncharged polymeric form of vanadic acid. The data in Table V indicated by superscripts b and c show that the amount of vanadic acid extracted by methyl isobutyl ketone decreased with time, although there was no effect on the distribution coefficient. With this solvent, also, there was no appreciable extraction below 5 M hydrochloric acid.
INDUSTRIAL AND ENGINEERING CHEMISTRY
July 1951
T~~~~VII,
ratio of chromium t o vanadium was varied as well as the concentration of each. The data of Table VI1 show t h a t the method is accurate t o approximately 1t o 2 parts per 1000. KETONEPHASE.The ketone phase samples were water extracted and the successive water washings pooled for analysis. The samples destined for polarographic analysis were reduced t o a suitable volume by evaporation.
vERIFICATIoN OF A~~~~~~~~~ M~~~~~ FOR H~~~
CHROMIUM-LOW VANADIUM MIXTURES
Chromium Content, Millimoles/Liter
Vanadium Content, Millimqles/Liter
Vanadium Found, Millimoles/Liter
25
0,1666
0
0.1666
25
0.833
0
0.833
25
1.666
0.1662 0,1664 0.1664 0.1664 0.831 0.834 0.832 0.830 1.666 1.665
0
1.666
25
2.500
0
2.500
50
0.417
0 5
Results. The data of Table VI11 show the effect of varying the sodium dichromate concentration while the vanadic acid concentration was maintained constant. A limited range of acid concentrations was investigated, for except a t 5 M hydrochloric acid and in one case at 4 M hydrochloric acid, no detectable amount of vanadium was found in the ketone phase. The minimum amount detectable by the polarograph was assumed t o be present, this being the most conservative figure. Figure 10 shows the variation with initial aqueous hydrochloric acid concentration of the concentration of vanadic acid in the aqueous phase a t eouilibrium. I n view of the excellent separation obtainable, an e x p e r i m e n t presented in Table I X was performed t o test its applicability t o the quantitative a n a l y s i s of c h r o m i u m a n d vanadium mixtures.
1.667 1.665 2.502 2.497 2.504 2.504 0.416 0.416 0.415 0.418 3.330 3.325 3.332 3.329
0.417 3.333
0
3.333
0
0.0417a
1683
0.0412a 0.04W
6 These results represent the smallest amount of vanadium determinable b y the polarograph.
OF SODIUM DICHROMATE AND VANADIC ACIDBETWEEN METHYL TABLE VIII. DISTRIBUTION ISOBUTYL KETONEAND WATERIN THE PRESENCE OF HYDROCHLORIC ACID
~Original ~ HC1
2.0 3.0
4.0
5.0
Initial Concn.,
T ~l ~ ~ Millimoles/Liter ~. , ~ i C.
NazCrtOi
HVOa
25 25 25 25 0 25 25 26 0 0 25 25 25 0
20 100 20 100 100 20 100 300 100 300 20 100 300 100 300
100 100 100 100 100 100
0
100
100 100 100 I00 100 100 100 100
Equilibrium Concn., Millimoles/Liter Aqueous Phase ~ ~ Ketone Phase NazCrzO? HVOs NazCrzOr HVOa
@a
..... .....
...
...
.. .. .. ..
.....
0.583 4.023 0.924 0.3010 0,868 15.21 0.762 4.66 0.2050 0.595 7.47 0.2167 2.627
96.8 100.9 99.5 89.2 97.4 107.6 96.5 108.7 75.6 82.4 102.4 86.3 103.4
17.10 87.8 96.2 16.36 88.9 256.8 96.8 269.2 16.09 85.1 251.0 96.0 274.4
2.5006 2.5006 2.5OOb 3.402 2.500b 2.500b 2.5006 2.5006 18.26 11.61 2.5006 8.40 2.500b
or 8, (theo.) equals K for *odium dichromate over K for vanadic acid. o vanadium detected; minimum amount detectable by the polarograph
SODIUM DICHROMATE-VANADIC ACID-HYDROCHLORIC ACIDWATER-METHYL ISOBUTYL KETONE
Analytical Methods. The analysis of either chromium or vanadium individually presents no problem as t o technique or accuracy; however, their mixtures present a difficult problem. The methods commonly used depend on the selective oxidation or reduction of either chromium or vanadium in a mixture of the two. No reagent t h a t is absolutely selective for this process has been reported. When both are present in considerable amounts, the error introduced by oxidizing or reducing a small portion of one with the other is usually insignificant. With solutions containing a large amount of one and a small amount of the other, this error is appreciable. Since it was anticipated that the ketone phase could contain a large amount of chromium with minute amounts of vanadium, a more accurate method was sou ht. T%e procedure developed is essentially the polarographic method for their analysis in steels reported by Lingane and Meites (16)with a few modifications. It is described in detail by Weinhstrdt and Hixson (66). Several standard solutions of vanadium and chromium mixtures were prepared from analyzed solutions of each. The
. ..
.....
1140 880 4150 1385 3980 730 4920 2500 , 324 1030 1378 4530 432
B(theo.)a
.
4430 2800 1950 1340 9530 737 557 214 3020 720 219 171 76.1 865 250
Figure 10. Effect of Hydrochloric Acid Concentration on Vanadic Acid Concentration in Aqueous Phase 4
Discussion. The separation factor, 8, was arbitrarily defined as the ratio of the distribution coefficient of dichromate t o t h a t of the vanadic acid at any given concentration. Since the distribution was assumed t o be present. coefficient is a function of the acid and salt concentrations, so also is the separation factor. As shown in Table VIII, the separation factors observed are greater than those predicted from the individual distribution coefficient values. This was apparently caused by a mutual salting-out effect. It was necessary t o base the calculation of several of the separation factors on the minimum amount of vanadium detectable by the polarograph. This accounts for the apparent decrease
TABLE IX. QUANTITATIVE ANALYSIS OF CHROMITJM A N D VANAEXTRACTION WITH METHYLISOBUTYL KETONEAT ROOM TEMPERATURE
DIUM BY
Molarity A ueoua %ci
Chromium +6 Content, Meq.
Vanadium +5 Content, Meq.
2
5.989
0.992
2
5.989
0.992
+
Chromium 4-6 Vanadium 5. Found, Found, Meq. Meq. Analysis by Extraction
5.984
0.986
Analysis by Polarbgraph
5.988
0.994
1684
I N D U S T R I A L A N D E N G I N E E R I N G CHEMISTRY
in the measured separation factor a t 3 M hydrochloric acid concentrations, In view of the minute amounts of vanadium found in the ketone phase a better indication of the excellence of the separa tion is the variation of the Concentration of vanadic acid in the aqueous phase a t equilibrium. Figure 10 shows t h a t the concentration of vanadic acid in the aqueous phase increaPes rapidly when the acid concentration is decreased from 5 to 4 M . Below 4 .LI hydrochloric acid there is relatively little change in the vanadic acid concentration. This
Vol. 43, No. 7
layers. The results obtained by extraction show excellent agreement with the known values. The results of a polarographic analysis of the mixture are also presented. Considering the simplicity of the method, further investigation t o determine its applicability to quantitative analysis should he made, CONCLUSIONS Vanadium and chromium may be separated by liquid-liquid extraction using methyl isobutyl ketone as the selective solvcnt. The optimum separation is obtained a t an aqueous hydrochloric acid concentration less than 3 M , a t 0 ” C., and with sodium dichromate concentrations in the range 0.02 to 0.3 M . With these conditions the reactions wit’h hydrochloric acid are (wentially eliminated and the separat,ion factor is in ewess of 4000. NOMENCLATURE P = separation factor, K for Xa&rzO,/K for HV03 C = concentration in millimoles/liter K = distribution coefficient, millimoles per liter, organic phase/ millimoles per liter, water phase LITERATCRE C I T E D (1) Axelrod, J., and Swift, E. H., J . Am. Chevi. Soc., 62, 33 (1940). (2) Dodson, R. W.,Forney, G. J., and Swift, E. H., I b i d . , 58, 2573
F i g u r e 11. Effect of Hydrochloric Acid C o n c e n t r a t i o n on D i s t r i b u t i o n Coefficient
indicates that a shaip increase in vanadic acid takes place between 4 and 5 Ill‘ hydrochloric acid. This mas observed in the study of the vanadic acid system, although it is not obvious in Figures 7 and 8. To emphasize this effect, the distribution coefficient is plotted in Figure 11 against initial aqueous hydrochloric acid concentrations for a 100-millimole initial vanadic acid concentration at 25 O C. On]?- one representative concentration is shown siiice the entire series exhibits the same trend. The curve for 100-millimole initial sodium dichromate concentration is for comparison. The sharp increase between 4 and 5 M hydrochloric acid is readily apparent. While the distribution coefficient does vary between hydrochloric acid concentrations of 1 to 4 d l , its value is extremely small; therefore, the vanadic acid ivould favor the aqueous phase almost exclusively. Any change in hydrochloric acid concentration in this iange would produce little effect on the aqueous vanadic acid concentration. At the high ratio of dichromate to vanadic acid (300/100) the data in Table VI1 show that the concentration of vanadic acid found in the aqueous phase a t equilibrium is actually larger than the original solution. This was caused by a decrease in volume of the aqueous phnse when mixed with the ketone. Since the vanadic acid 71 as almost exclusively in the aqueous phase the concentration became greater than the original. Table IX presents results t h a t further emphasize the excellence of the separation. These data were obtained by successively extracting with methyl isobutyl ketone a 2 M hydrochloric acid solution containing known amounts of vanadic acid and sodium dichromate. The chromium reported was in the ketone extracts, while the vanadium shown was entirely in the aqueous raffinate
(1936). (3) Doeuvre, M. J., Bull. soc. chim., 39, 1597 (1926). (4) Eppley, M., and Vosburgh, W. C., J . Am. Chem. Soc., 44, 2148 (1922). (5) Forbes, G. S., and Bartlett, E. P., Ibid., 35, 1535 (1913). (6) Foster, M. D., U. S. Geol. Survey, B d l . 950, 15-18 (1946). 41, 2298 (7) Garwin, L., and Hixson, A . N., ISD. Ems. CHEM., (1949). (8) Giahame, D. C., and Seaborg, G. T., J . Am. Chevi. SOC.,60,2624 (1938). (9) Hixson, A. W.. and Miller, R., U. S. Patent 2,227,833 (Jan. 7, 1941). (10) International Critical Tables, National Auademy of Sciences, Vol. I, p. 202, New York, McGraw-Hill Hook Co., 1926. (11) Jilek, A., and Vicovsky, IT,, Collection C’eechoslov. Chena. Con+ mun., 4, 1 (1932). (12) Kayto, S., and Isii, R., Sei. Papers Inst. P h y s . Chem. Rcserirch ( T o k g o ) , 36, 82 (1939). (13) Klein, I. J., Ph.D. dissertation, Columbia Vniversity, 1942. (14) Langmuir, A. C., J . Am. Chem. Soc., 22, 102 (1900). (15) Lingane, J. J., and hleites, L., Jr.. A s . 4 ~ .CHEM.,19, 159 (1947). (16) Lingane, J. J., and Meites, L., Jr., J . a m . Chem. Soc., 68,2443 (1946). (17) McNabb, W.hI., private communications. (18) hfellor, H. W.,“A Comprehensive Treatise on Inorganic and Theoretical Chemistry,” Vol. XI, p. 244, New York. Longmans, Green & Co., 1931. (19) Montequi, R., and Gallego, hI., Anales soc. espafi. fk.q z ~ i r n .32, , 134 (1934). (20) Mylius, F., Ber., 44, 1515 (1911). (21) Mylius, F., 2. anorg. Cham., 7 0 , 205 (1911). (22) Sachtrieb, N. H., and Conway, J. G., J . A m . Chem. Soc., 70, 3647 (1948). (23) Nachtrieb, N. H., and Conway, J. G . , “Studies in Complex Formation. I. The Extraction of Feriic Chloride by Isopropyl Ether,” declassified papers, Dept. of Commerce, Kashington, D. C. (24) Noyes, A. A,, J . Am. Chem. Soc., 30, 51.5 (1908). (25) I b i d . , p. 558. (26) Rice, F. O., and Fryling, C. F., I b i d . , 47, 379 (1925). ( 2 7 ) Rothe, J. W., Chem. News, 66, 182 (1892). (28) Rothschild, B. F., Templeton, C. C., and Hall, N.F., J . Phus. & Colloid Chem., 52, 1006 (1948). (29) Sandell, E. B.. I s n . Eso. CHEM.,ASAL. ED., 8, 336 (1936). (30) Serfass, E. J., I b i d . , 12, 636 (1940). (31) Skey, IT., Chem. ,\’ews, 36, 48 (1877). (32) Speller, F. N., Ibid., 85, 124 (1901). (33) Swift, E. H., 6.Am. Chem. Soc., 46, 2375 (1924). (34) Vilbrandt. F. C., and McCormack, R., Bd2. J‘iroinin Polgtccir. I n s t . . Eng. Ezpt. Sta. H u l l . Ser. No. 64, 10 pp. (1946). (35) Weinhardt, A. E., and Hixson, 9.S . , Ph.D. dissertation, University of Pennsylvania, 1949. (36) Wells, J. F., and Hunter, D. P., A n a l y s t , 73, 671 (1948). RECEIVED June
13, 1950
