790
M.
CEFOLA4,
R.
c. TAYLOR, P.s. GEKTILE,A S D A. TT. CELIANO
Yol. 66
COORDINATION COMPOUNDS. 111. CHELATE CONPOUNDS OF THE URANYL ION WITH HYDROXY, MERCAPTO, AND AMINO ACIDS’ BYMICHAEL CEFOLA,ROBERT C. TAYLOR, PHILIPS. GENTILE,AND A. V. CELIANO Department of Chemistry, Fordham University, New York, N . Y. Received July
17,1961
This investigation involved a potentiometric study of the chelates formed by the uranyl ion with the amino, hydroxy, and mercapto analogs of acetic, propionic, and succinic acids. The order of decreasing stability was aspartic, p-alanine, glycine, malic, thiomalic, p-hydroxypropionic, glycolic, and thioglycolic acids, The stability of chelates with analogs of acetic, propionic and succinic acids decreases for any given acid as the donor group changes from XH2 to O H to S H . This decrease parallels the order of decreasing basicity in any given series. Considering acids which contain the same donor group (“2 or OH or S H ) decreasing stability is observed in the analogs of succinic, propionic, and acetic acids, in that order.
Titration studies of the uranyl ion with acids and complexing agents has been virtually neglected except for a few ligands.2-8 This investigation was undertaken to establish the stabilities of the uranyl chelates of several hydroxy, mercapto, and amino acids, and to determine trends in their stabilities. The ligands were chosen so as to be able to determine the effect of ring size, ligand basicity, and nature of the donor atom on the stabilities. A polarographic study of these compounds, as well as of the uranyl chelates of p-diketones and several types of organic sulfur compounds, also was undertaken and will be reported at a later date. Little work has been done with compounds containing sulfur coordinating groups because of their tendency t o oxidize to disulfides, especially in the presence of metal ions. Experimental Materials.-All chemicals used were either reagent grade or purified. Organic acids were prepared at a concentration of 0.1 M and were diluted for each titration so that the final solution contained about 1 X 10-3 mole of acid. p-Hydroxypropionic acid was prepared by dissolving @-propiolactonein hot water, to convert the lactone to the acid.g The acids, with the exception of glycine and p-alanine, were standardized by potentiometric titration with potassium hydroxide. Glycine and @-alanine were standardized with potassium hydroxide in the presence of forrnaldehyde.’O The uranyl nitrate solution was prepared at a concentration of 0.033 M and diluted so that the final solution contained between 1 X 10-8 and 1 X 10-5 mole of metal ion. It was standardized by the addition of hydrogen peroxide and the titration of.the nitric acid formed by the resulting reaction.ll Potassium hydroxide was prepared carbonate-free a t a concentration of 0.1 M by the method of Schwarzenbach and Biedermann’2 and standardized with potassium acid phthalate. Apparatus and Procedure.-The experimental method consisted of the potentiometric titration of the arganic acids in the absence and in the presence of uranyl ion. The acids employed were glycolic acid, thioglycolic acid, glycine, (1) This study was supported by a grant from the U. S. Atomic Energy Commission, AT-(30-1)906. (2) ill. Cabell, Analyst, 7 7 , 859 (1952). (3) R. Irving and H. Rossotti, J . Chew&.Soc., 2910 (1954). (4) 8. Ahrland, Acta Chem. Scand., 6 , 199 (1951). ( 5 ) 9. Ahrland, ibid., 3, 783 (1949). ( 6 ) S. Ahrland, %bid.,7 , 485 (1953). (7) J. Rodgers and ITr. Newman, Stomic Energy Commission Report 4ECO-2218, August 11, 1948. (8) R. M. Isatt, W. C. Fernelius, and B. P. Block, J Phys. Chem., 69, 80 (1955). (9) T. L. Gresham, J. E. Jansen, and W. Shaver, J. A m . Chem. Soc., 70, 998 (1948). (10) I. Kolthoff and N. Furman, “Volumetric Analysis,” John Wiley and Sons, Ino., New York, N. Y., 1929, p. 164. (11) W. Bunce, G. Morrison, W. Chorney, and R. Nundy, Atomic Energy Commission Report ‘No. 2-1044, July 7, 1944. (12) G. Sohwarsenbaoh and W.Biedermann, Helw. Chim. Acta, 31, 339 (1948).
malic acid, thiomalic acid, aspartic acid, p-hydroxypropionic acid, and p-alanine. The ionic strength was maintained at, 0.1 by using 0.1 IVT potassium chloride as a supporting electrolyte. All measurements were carried out in an atmosphere of porepurified nitrogen a t a constant temperature of 30.0 k O . 1 . Titrations were carried out using 1:1, 2:1, and 10: 1 ratios for the acid to metal ion concentration. The p H values were determined with a Beckman Model G p H meter with extension glass and calomel electrodes. Interpretation of the data is simplified if hydrolysis, which is p H dependent, may be neglected. It is possible by varying metal and ligand concentration to find a region of negligible hydrolysis. The agreement of the values of stability constants obtained a t the different metal ion concentrations indicates that the amount of hydrolysis was negligible in the region between p H values of 2 and 4. It has been e~tablished’~J~ that UOz++ is the species of hexavalent uranium existing in solution of pI-1 approximately equal to 3 . Due t o early precipitation of the metal hydroxide in the 1O:l runs for glycine and @-alanine,calculations could not be made. Formation function curves did not permit calculation of KZ for uranyl aspartate. The values reported are the averages of several runs. Acid dissociation constants were determined potentiometrically and agree well with previously reported values.16-1g The logarithms of the stability constants of the uranyl chelates are listed in Table I and were calculated by Bjerrum’s graphical method.20 The error limitszf are: f 5 % for K,, & l o % for KZ. This corresponds to the errors: 1 0 . 0 2 for log K I , f 0 . 0 5 for log KP.
Discussion The over-all order of decreasing stability for the ligands was aspartic, p-alanine, glycine, malic, TABLE I STABILITY CONSTANTS OF MONOBASIC AND DIBASIC ACIDS WITH URANYL YITRATE Acid
log K1
log Ki
Aspartic @-Alanine Glycine Malic Thiomalic p-Hydroxypropionic Glycolic Thioglycolic
8.00 7.78 7 53
7 53 7.15
5.50 3 56 3 25
2 97 2 88
.. 3.63 3 42
2.88 2 40 2 40
(13) S. Ahrland, Acta Chem. Scand., 3 , 374 (1949). (14) M. Crandall, J. Chem. Phys., 17, 602 (1949). (15) R. Canaan and A. Kibrick, J . Am. Chem. Soc., 60, 2314
(1938). (16) N. C. Li and R. A. Manning, ibid., 77, 5225 (1955). (17) H. Kroll, ibid., 74, 2034 (1952). (18) R. Taylor and M. Cefola, Doctoral Dissertation, Fordham University, June, 1957. (19) N. C. Li and E. Doody, J. A m . Chem. Soc., 7 2 , 1891 (1950). (20) J. Bjerrum, “Metal Ammine Formation in Aqueous Solution,” P. Haase and Son, Copenhagen, 1941. (21) G. R. Choppin, private communication.
HEATSOF COMBCSTIOS OF Resa ASD Re&
May, 1962
thiomalic, P-hydroxypropionic, glycolic, and thioglycolic acids, as shown in Table I. The stability (of chelates with analogs of acetic, propionic, and succinic acids decreases for any given acid as tho donor group changes fromNHs t o OH to SH. This decrease parallels the order of decreasing basicity in any given series. Considering acids which contain the same donor group (NH2 or OH or 8H) decreasing stability is observed in the analogs of succinic, propionic, and acetic acids, in that order. This order is found to be unusual in that the six-member ring chelates of the propionic acid analogs are more stable than the five-member ring chelates of the acetic acid analogs. Succinic acid analogs, although the weakest bases, form the strongest chelates in any group. These would be expected to be the least stable inasmuch as a decrease in ligand basicity results in a corresponding lower affinity for metal ions. The behavior of succinic acid analogs may be attributed to the formation of either a bidentate or tridentate complex. The bidentate complexes possibly may possess either a five- or six-membered ring, thus taking on the structure of acetic acid analogs with a mbstituted carboxylmethyl group, or of propionic acid analogs with a substituted carboxyl group. I n either case, the bidentate chelates would be expected to be only slightly stronger than tbose of the acetic acid or propionic acid analogs, whereas tridentate formation would be expected to give a large increase in the stability constant (-2 log units).22 Such an increase is observed only for malic scid. Lumb and Martel123 ( 2 2 ) M. Cefola, A. Tompa, A. InoTg. Chem., 1, 290 (1962).
V. Celnano, and P. S. Gentile,
791
have observed that aspartic acid forms bidentates with alkaline earth ions and for these chelates the stability constant was 0.2log unit higher than the glycine, the small increase being attributed to the inductive effect of the negative carboxylate group in aspartic mid, which can lend stability to the structure by increasing the basicity of the donor group toward the metal ion. The difference between the stabilities of the aspartate and p-alinate chelates of uranyl ion is approximately 0.2 log unit, thus indicating formation of a bidentate chelate with aspartic acid. I n all probability aspartic acid forms a six-membered ring since the p-alinate is more stable than the glycinate. Table I1 lists the relative position of U02t2 in a series of metal ions for the formation of glycinates and aspartates. TABLB I1
STABILITY CONSTANTB OF VARIOUSMETALSWITH GLYCINE AND
-Glycine---Metal ion
Cu(I1) UOZ Xi( 11) Zn(I1) Co(I1) Pb(I1) Mn( 11) Ag( 11)
log
ASPARTICACID
Ki
8.62(24) 7.53 6.18 (24) 5.52(24) 5.23(24) 5.47 (24) 3 44 (24) 3.51 (24) I
-Aspartic AMetal ion
Cu(I1)
uoz
Ni( 11) Co(I1) Zn(I1) Cd(I1) Hg(1I) Ca( 11)
acidlog Kt
8.57(25) 8.00 7.12 (25) 5.90(25) 5.64(25) 4.37 (25) 2.43 (23) 1.60 (23)
(23) R. F. Lumb and A. E. Martell, J . Phys. Chem., 87, 690 (1953). (24) G. Monk, TTans. Faraadag Soe., 47, 297 (1951). (25) S. Chaberek and A. E. Martell, J . Am. Chem. Soc., 74, 6021 (1952).
THE HEATS OF COMBUSTION OF Resz AND RezS7 AND THE THERMODYNAMIC FUNCTIOKS FOR TRASSITION METAL SULFIDES'~2 BYJ. E. MCDONALD~ AND J.
w.COBBLE
Department of Chemistry, Purdue University, Lufayette, Indiana Received Julv Zh. 1861
The heats of combustion of Re& and Red37 have been measured by bomb calorimetry. The heat and free energy of formation a t 25' of Re& are -42.7 f 1.2 kcal. mole-l and -41.5 =!= 1.2 kcal. rnole-l, res ectively. The similar quantities for Re& are - 107.9 d: 1.8 kcal. mole-' and - 101.0 f 1.8 keal. mole-', respectively. &ith data on these key sulfides, along with recalculated heate of formation from previous vapor pressure measurements, it has been possible to estimate the thermodynamic functions for a number of sulfides not previously available.
I. Introduction Very few thermodynamic data have been available on transition element sulfides in spite of their potential interest in solid state and high temperature chemistry. Certain vapor pressure meas(1) This research was supported by the United States Air Force through the Air Force Office of Scientific Research of the Air Research and Development Command under contract AB 18(600)-1525. Reproduction in whole or part is permitted for any purpose of the United States Government. (2) This constitutes communication number VI11 in our previous series on rhenium and technetium chemistry. For the previous paper in this series see J. A m . Chem. Soc., 82, 2111 (1960). (3) From the Ph.D. thesis of J. E. McDonald, Purdue University, 1961; Dow Chemical Co. Fellow, 1959-1960; Procter and Gamble Fellow, 1960-1961.
urements of Biltz, Juza, and their co-workers fix the heats of formation of Pt&, PtS,4 IrzS8, ReSz,s Rust,' and OSSZ~ if the assumption 18 made that the only vaporizing species is gaseous sulfur. Since Re& can be directly sublimed at -lOOOo, this assumption needs independent verification. The purpose of this communication is to report the calorimetric determination of the thermody(4) W. Bilta and R. Juza, 2.anorg. Chem., 190, 161 (1930). (5) W. Biltz, J. Laar, P. Ehrlich, and K. Meisel, zbid., 283, 257 (1937). (6) R. Juza and W. Biltz, 2. Elektrochem., 37, 498 (1931). (7) R. Juza and W.Meyer, 2. anorg. Chem., 213, 273 (1933). (8) R. Juaa, ibiad., ZlS, 129 (1934).
