Aug., 1956 sodium hydroxide to nickelous chloride solution, followed by washing by decantation until the black gel was free of chloride ions; (D) nickel oxide prepared by heat-treatment of Ni(OH)2 a t 1100” in a current of nitrogen for 2 hours (“active” oxygen content ca. 0.02% and crystal size about 1500 A,); (E) nickel oxide as in (D) except for heat treatment a t 300” (“active” oxygen content 0.37%, crystal size 104 and (F) Ni(OH)2 prepared by precipitation from nickelous nitrate solution, followed by washing by decantation until the green gel was free of nitrate ions, and drying at room temperature. Presumably the significant characteristic of the edge is the location of the first principal peak, since this is thought to be the result of 1s-4p transitions. In Fe and Mn this feature of the trivalent oxides is displaced some 3 or 4 volts toward higher energies than in the respective divalent oxides. Figure 1 shows the K edge structure for various nickel oxides. The zero shown is taken as the top of the Fermi band in pure n i ~ k e l . ~ The six curves are labeled as indicated in the previous paragraph. Curves A and B are very similar, there being no significant difference between them with regard to position or shape. Curves C to F are likewise very much alike, but there are small, yet significant, differences which deserve comment. The principal peak for samples A, D and E occurs at the same position. Our measurements place the peak for sample C at about 0.75 volt higher, whereas sample F [Ni(OH)2]peaks about 0.75 volt lower. Inasmuch as these data are obtained by simultaneous runs, the shifts are real and are considered t o be precise to about 0.25 volt. However, the shifts are far smaller than one finds in the corresponding Fe and Mn compounds. It is noteworthy that there is a virtual absence of low energy absorption in sample C. The oxides of all of the previously examined elements of the first transition series show such an absorption, and hence the results for the black, nickel oxide (sample C) represent an anomaly. It may be more than coincidental that samples D and E, containing small amounts of active oxygen, likewise exhibit a comparatively small low energy absorption, and are black in color as is sample C. It is this lack of low energy absorption which apparently accounts for the results of Cairns and Ott. Since they measured the shift in the edge by photographic methods, thick samples of nickel oxide would give considerable absorption before the principal peak was reached. Although it is shown that the Cairns-Ott measurements are not definitive, the present results do not constitute proof that higher nickel oxides do not exist, but suggest that if the trivalent or other higher oxides of nickel exist, their K edge structure does not bear the same relation to the edges of the divalent oxides as do the edges in the corresponding oxides of iron and manganese which have been measured. This is a likely possibility inasmuch as lattice symmetry and the homopolar character of the bonding may be involved in determining the exact edge structure.
NOTES
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w.);
(9) W. W. Beeman and H. Friedman, Phyo. Rsv., 46,392 (1939).
I
0
IO
I
I
1
I
20
30
40
50
ELECTRON VOLTS, Fig. 1.-K absorption edges of various oxides of nickel. The zero is chosen as the top of the Fermi band in pure nickel.
One of us (H. P. H.) gratefully acknowledges financial assistance from The Robert A. Welch Foundation. T H E SOLUBILITIES OF SOME METAL NITRATE SALTS I N TRI-n-BUTYL PHOSPHATE BY WESLEYW. WENDLANDT AND JOHN M. BRYANT Department of Chemistry and Chemical Engineering, Texas Technological College, Lubbock, Texaa Received February 90, 1966
The use of tri-n-butyl phosphate (TBP) as a solvent in the liquid-liquid extraction of metal ions from aqueous solutions is well known. The rare earths, thorium,2 uraniuma and plutonium4 have all been extracted from aqueous solutions using ( 1 ) D. F. Peppard, J. P. Faris, P. R . Gray and G. W. Mason. THIS JOURNAL, 57, 294 (1953); B. Weaver, F. A. Kappelmann and A. C.
Topp, J . A m . Chem. Soc., 75,3942 (1953); A. C. Topp and B. Weaver, U. 8. Atomic Energy Commission, Rept. ORNL-1811, Oat. 15, 1954; J. G . Cuninghame, P. Scargill and H. H. Willis, Atomic Energy Research Establishment, Harwell, Rept. AERE-C/M-215, Aug. 13, 1954; J. C. Warf, J . Am. Chem. SOC.,71,3257 (1949). (2) M. W. Lerner and G. J. Petretic, Anal. Cham., 88, 227 (1956); M. R. Anderson, U. S. Atomic Energy Commission, Rept. ISC-118, Dec., 1953. (3) H. T. Hahn, ref. 2, HW-32626. July 20, 1954. (4) B. Goldschmidt, P. Regneut and T. Prevot, Rapp. cant. el. nucl. Saclay, number 397, 1955.
1146
NOTES
this solvent under various conditions of acidity and foreign electrolyte concentrations. Analytical determinations have been developed for the separation of uranium,6 copper,G iron’ and aluminum* from aqueous solutions using TBP. To investigate other possible metal ions which may be separated by liquid-liquid extraction, a systematic study of the Solubilities of a series of metal nitrate salts was made in the pure solvent. Experimental Procedure Chemicals.-Tri-n-butyl phosphate was obtained from Eastman Organic Chemicals, Inc., Rochester 3, N. Y., and used without further purification. All the other chemicals were of reagent grade quality. The procedure for the solubility determinations consisted of adding about 25 g. of the solid salt to 20 ml. of T B P contained in a 50-ml. screw cap bottle, sealing the bottle,.and equilibrating the contents on a mechanical “wrist action:’ shaker for 48 to 72 hr. at room temperature, 25 to 27 It was found that this time was sufficient for equilibrium conditions to be established. At the end of this time, three phases were present in the bottles: a solid hydrated salt phase, an aqueous phase containing a saturated solution of the metal nitrate salt and an organic phase containing the dissolved metal salt. The organic phase was separated, centrifuged and analyzed for the metal salt content. The analysis consisted of weighing out 1 to 4 g. duplicate samples of the centrifuged organic phase into 125-ml. separatory funnels containing 25 ml. of benzene and 50 ml. of water. After an equilibration time of two minutes, the aqueous phase was removed, 50 ml. of water added and the equilibration repeated. Two such extractions were sufficient to remove the metal salt from the organic benzene phase. The metal ion contents in the extracted aqueous phases were determined by standard procedures .g
.
Vol. 60
TABLE I THE SOLUBILITIESOF METAL NITRATE SALTS IN TRI%-BUTYL P H O S P H A m
SaIt
UOz(NOs)z~6HzO Th(NOd)n.4Hz0 Fe(NOs)a.9Hz0 Bi( N03)s.5Hz0 La( NO&6Hz0 CU(N03)z.6Ht0 Cd(N03)~.4HzO Zn( NOs)2.6H20 Hg(NOs)z Ca(NOa)2.6H~0 Co(N03)z*6HzO LiN03.3Hz0 Mg(NOa)z*6H20 Ni(NO&.BHzO Al( NOs)a*9Hz0 AgNOa Sr(N0a)z Pb(NOa)z Ba(NOd2 Solubilities expressed in g. of of solution. (1
Solubilityo g./100 g. soln.
43.6 42.6 34.8 28.6 28.5 21.5 21.5 20.6 18.5 17.0 13.5 12.0 11.6 10.4 9.37 2.59 0.81 0.39 0.00 anhydrous salt
43.4 42.4 33.9 28.6 28.3 21.3 21.4 20.6 19.0 17.0 13.4 11.9 11.4 10.5 9.45 2.59 0.81 0.38 0.00 per 100 g.
aly occurs in that magnesium nitrate is less soluble than calcium nitrate. The order of solubilities of the remaining alkaline earth nitrates decreases sharply. The possibility of extracting other metal ions from solution by liquid-liquid procedures with TBP seems to be quite favorable. Under certain conditions, it should be possible to extract about the first 15 metal salts in the table. Further work is being conducted on these separation possibilities.
Results The solubilities of the metal nitrates are given in order of decreasing solubility in Table I. The most soluble salts are those of uranium(V1) and thorium. Thus, it can be seen why liquid-liquid extraction procedures are so effective with these ACETONITRILE-WATER LIQUID-VAPOR two metal salts from nitric acid solutions. EQUILIBRIUM The solubility of the metal nitrates is due to the formation of molecular addition complexes between BY F. D. MASLAN AND E. A. STODDARD, JR. the salt and the TBP. This complex for uranium Petrochemicals Department, National Research Corporation, Cambridge, (VI) nitrate has the formula [UOz(TBP)~(NO&].lo Mass. Cerium(1V) nitrate also forms such an addition Received Febrzlarg PO, 1966 complex but having the composition [Ce (TBP)z I n the course of a research project a t this LaboThe extraction of the other metal niratory, the distillation of acetonitrile-water mixtrates obviously forms molecular addition complexes but as yet they have not been investigated. tures has received attention. A preliminary exThe solubility of lanthanum nitrate is very periment indicated a difference from the azeotrope similar to the transition metal nitrates. The fairly a t 760 mm. as published by Othmer and Josefogreat solubility of this representative of the rare witz.l Further examination of the literature indiearth group explains why separation procedures cated an uncertainty as to the composition and boiling points of the azeotrope. Othmer, et al., have been so successful using TBP. The solubilities of metal nitrates in any one give 72.60 mole % acetonitrile at the azeotrope, group decrease with increasing atomic weight. while an industrial source gives 69.2.2 Our prelimiHowever, with the alkaline earth group, an nnom- nary results indicated a figure closer to the latter. I n order to clear u p this matter, the liquid-vapor (5) W. B. Wright, U. 9. Atomic Energy Commission Rept. Y-884, equilibrium for acetonitrile-water at 1 atmosphere Oct. 18, 1954; R. J. Guest, Dept. of Minea and Technical Surveys, was determined. Canada, Rept. NP-5763, May 30, 1955. ( G ) L. M. Melnick and H. Freiser, Anal. Chem., Z7, 462 (1955). (7) L. M. Melnick and H. Freiser, %bid.,25, 856 (1953). ( 8 ) M. Aven and H. Freiser, Anal. Chim. Acta, 6 , 412 (1952). W. Scott, “Standard Methods of Chemical Analysis,” (9) Vol. I, edited by N. H. Furman, D. Van Nostrand Co., Inc., New York, N. Y . , 5th edition, 1946; G. E. F. Lundell, H. A. Bright and
Materials.-The acetonitrile for this work was purified as follows. Eastman research grade acetonitrile was distilled in a aO-plate, 1-inch diameter, perforated plate column. A reflux ratio of 10 t o 1 was used. The middle cut, boiling a t 82’, was retained. Carbide and Carbon reports 81.8’
J. I. Hoffman, “Applied Inorganic Chemistry,” John Wiley and Sons, Inc., New York, N . Y., 2nd edition, 1953. (IO) R. L. Moore, U. S. Atomic Energy Cornmission, Rept. AZCD3106, July 10, 1961.
(1) D. F. Othmer and 9. Josefowitz, Ind. Eng. Chem., 89, 1175 (1947). (2) Carbide and Carbon Chemicals Co., “Acetonitrile,” February, 1955.
W.
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