NOTES
Aggregation Studies of Some Nuclear Magnetic Resonance Shift Reagents by Vapor Phase Osmometry Jean F. Desreux,‘ Lloyd E. FOX,^ and Charles N. Reilley Department of Chemistry, University of North Carolina, Chapel Hill,N . C . 27514
LANTHANIDE METAL COMPLEXES have been used extensively as shift reagents for simplifying NMR spectra since the first report by Hinckley ( I ) . Among these compounds, Eu(II1) and Pr(II1) 2,2,6,6-tetramethyl-3,5-heptanedionechelates (thd), as suggested by Sanders and Williams (2),and Eu(II1) and Pr(II1) 1,1,1,2,2,3,3-heptafluoro-7,7-dimethyl-4,6-octanedione chelates (fod), as suggested by Rondeau and Sievers ( 3 ) , appear to be of the greatest interest for numerous spectrometric applications. The molecular associations of these chelates in solution have not been thoroughly investigated although their dimerization in the solid state has been reported (4,5), and potential existence of aggregates in solution phase has been suggested (3). Characterization of the molecular states of these species in the absence of a substrate must be a first step in the accurate interpretation of isotropic shifts. In the past, it has been commonly assumed that the only components in solution are the free organic molecules under study and the adduct with one molecule of chelate. If any chelate oligomer is present, such a simple model would no longer be valid. Many of the lanthanide-thd chelates are aggregated (4, 5 ) in the solid state. The lighter lanthanides, La(thd)a to Sm(thd)3, crystallize from hexane exposed to the atmosphere as anhydrous dimers and, as the ionic radius is decreased, the monohydrate also starts to crystallize. Dy(thd)r is obtained only in the form of a hydrated dimer; hydrated monomers are obtained for the heavier lanthanides [above Ho(thd),]. The present study was, therefore, undertaken to investigate by vapor phase osmometry the aggregation of the thd and fod chelates of Eu(II1) and Pr(II1) in a few typical solvents, chloroform, carbon tetrachloride, and hexane. EXPERIMENTAL
The osmometric measurements were carried out at 37 “C with a Mechrolab-301 osmometer. A R ; , the thermistor resistance read from the osmometer, was extrapolated for a time t = 0 (ARO) as recommended by Meeks and Goldfarb (6) and by Kume and Kobayashi (7). No drop size effect was On leave from the University of Liege (Belgium). Present address, The Upjohn Company, Kalamazoo, Mich. (1) C. C. Hinckley, J. Amer. Chem. Soc., 91,5160 (1969). (2) J. K. M. Sanders and D. H. Williams, Chem. Commwz., 1970, 422. ( 3 ) R. E. Rondeau and R. E. Sievers, J. Amer. Chem. Soc., 93, 1524 (1971). (4) C.S. Erasmus and J. C . A. Boevens. Acta CrystaNoar., - . B26, 1843 (1970). (51 . , C. S. Erasmus and J. C. A. Boevens. J. Crvst. Mol. Struct.. 1. 83 (1971). (6) A. C. Meeks and I. J. Goldfarb, ANAL.CHEM.,39, 908 (1967). (7) S. Kume and H. Kobayashi, Makromol. Chem., 79,1 (1964). -
-
0
2
4
6
8
1
0
c x IO*(mote/liter)
1
2
Figure 1. Vapor phase osmometric behavior of Eu(II1) and Pr(II1) fod and thd chelates in chloroform referenced to benzil at 37 “C. The ordinate is proportional to the vapor pressure difference between the dry solution and the dry solvent 0 Benzil, A
Eu(fod)8,H Pr(fod)a, X Eu(thdh,
+
Pr(thd)r
observed as long as measurements were performed in a reasonably consistent manner. Benzil was used as the standard, and all the solvents were purified by the usual techniques (a), dried, and checked by gas chromatography. Chloroform had to be purified with special care because of an acidic impurity which decomposes the lanthanide chelates. All solutions were prepared in a dry atmosphere. The commercially available Eu(II1) and Pr(II1) thd chelates were impure as evidenced by abnormally low osmometric molecular weights. Also, the melting points were not ia agreement with the reported values of Sievers et al. (9) even after extended drying under vacuum. Sublimation at less than 0.1 mm Hg and 180 “C only slightly improved the purity as did successive recrystallizations from dry hexane. Purification by liquid chromatography on alumina and on silica gel columns failed since the complexes could not be
,
I
,
I
(8) J. A. Riddick and W. B. Bunger, in “Techniques of Chemistry,” Volume 2, “Organic Solvents,” A. Weissberger, Ed., WileyInterscience, New York, N.Y., 1970. (9) K. J. Eisentraut and R. E. Sievers, J. Amer. Chem. SOC.,87, 5254 (1965).
ANALYTICAL CHEMISTRY, VOL. 44, NO. 13, NOVEMBER 1972
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50
9
401
c x 102
(mote/liter)
Cxl
Figure 2. Vapor phase osmometric behavior of Eu(II1) and Pr(II1) fod and thd chelates in carbon tetrachloride referenced to benzil 0
Benzil, A E~(fod)~, Pr(f~d)~, X Eu(thd)r,
+
Pr(thd)a
eluted even with very polar solvents. Partial deactivation of the columns with dimethyldichlorosilane did not improve the elution although the technique was not exhaustively implemented. Noteworthy here is the possibility of recovering precious organic compounds following NMR measurements by separating them from the shift reagents by liquid chromatography. An alumina column should preferably be used since the weak acidity of the silica gel slightly decomposes the lanthanide complex liberating the chelating agent, H(thd), which is easily eluted. The synthesis of the thd chelate involves the reaction between a lanthanide nitrate and a sodium thd salt in an aqueous-alcoholic media as reported by Sievers et al. (9). By this method, one obtains compounds which give excellent osmometric and melting point data. The commercially available compounds can be purified by recrystallization from a 50z aqueous solution of ethanol. Nearly anhydrous crystals are obtained which were shown to be pure after drying and recrystallization from hexane. The purified compounds must be kept under vacuum and in a dry atmosphere. The osmometric behavior of the commercially available fod salts was not modified by the purification method described above. The purity of the thd chelates has a marked influence on the NMR shift. A 0.2-ppm increase in the shift of the Q CH?peak ofn-pentanol was observed with purified Eu(thd)s compared to the commercially available product for a molar ratio of shift reagent to alcohol of 0.2 in carbon tetrachloride. Further, several additional peaks have been observed in the NMR spectra of impure thd chelates. RESULTS AND DISCUSSION
The results obtained by osmometric measurements on solutions of Eu(II1) and Pr(II1) thd and fod chelates in chloroform, carbon tetrachloride, and n-hexane are summarized in Figures 1 , 2 , and 3. ARO, the thermistor resistance extrapolated to a time r = 0, is proportional to the vapor pressure difference between the dry solution and the dry solvent. The ratio of this value for the standard to the corresponding value 2218
Figure 3. Vapor phase osmometric behavior of Eu(II1) and Pr(II1) fod and thd chelates in n-hexane referenced to b e n d 0
Benzil, A Eu(fodh,
+
Pr(fod),, X Eu(thd)8,
Pr(thd)r
for the chelate at the same concentration is the mean aggregation number. Both complexes, Eu(thd), and Pr(thd),, were monomeric to the limit of experimental errors in the three solvents used and over the concentration ranges investigated. The Eu(fod), and Pr(fod), chelates form aggregates whose concentrations increase in the order chloroform, carbon tetrachloride, n-hexane. One may assume that the more polar nature of the fod chelate, due to the nonsyrnrnetrical nature of the P-diketone itself and to its fluorocarbon chain, partially account for the aggregation of these chelates. The smaller steric effect of the fluorocarbon chain could also be of importance. The osmometric data were interpreted with a computer program in terms of molecular associations via the Bjerrum relationship. Kertes et al. (10) proposed a very elegant way to solve this relationship without assuming a value for a concentration of the monomeric species. The value of S, in Equation 15 of Kertes' work (IO), can be varied by a factor of ten without any modification of the stability constants of of the aggregates. A linear least-squares procedure proposed by Sullivan, Rydberg, and Miller (11) was applied to Equation 21 of reference (10). A program which makes use of these two methods is particularly suitable for a small computer. The results of the best mathematical fit to the experimental measurements of each system are summarized in Table I. The aggregation sequence is in agreement with the well known solvation powers of these solvents. Chloroform, by its acidity, very likely interacts with the chelate. Armitage and Hall (12) recently reported that isotropic shifts of carbohydrates by Eu(thd)s are some two times smaller in this solvent than in carbon tetrachloride. In addition, (10) A. S. Kertes and G. Markovits, J. Phys. Chern., 72, 4202 ( 1968). (11) J. C. Sullivan, J. Rydberg, and W. F. Miller, Acta Clzern. Scund., 13, 2023 (1959). (12) I. Armitage and L. D. Hall, Can. J. Chern., 49, 2770 (1971).
ANALYTICAL CHEMISTRY, VOL. 44, NO. 13, NOVEMBER 1972
the solubility of Yb(thd)3 is much greater in chloroform than in carbon tetrachloride and in carbon tetrachloride than in n-hexane as expected. The P r ( f ~ d )complex ~ is systematically more aggregated This is in agreement with some than the E ~ ( f o d )complex. ~ related observations contained in a report by Sievers et a [ . (13) on the solvation behavior of 3-trifluoroacetyl-d-camphorate chelates. In chloroform, the La(II1) complex is much more aggregated than the Er(II1) complex. The decrease in aggregation with decreasing ionic radius may be interpreted through a lowering of the dipole moment as previously suggested for chromatographic data (14). Crystallographic data also show this influence of the ionic radius. Further, Yb(thd)3 is approximately ten times more soluble in chloroform than is Eu(thd), (15, 16), as can be expected from the smaller ionic radius of the Yb(II1) ion and, thus, from its smaller dipole moment and molecular association. On the basis of NMR spectra, Archer, Fell, and Jotham (17) recently reported the aggregation of the thd chelates. The relatively intense peak attributed by these authors to di(13) B. Feibush, M. F. Richardson, R. E. Sievers, and C. S . Springer, Jr., J. Amer. Chem. Soc., in press. (14) C. S . Springer, Jr., D. W. Meek, and R. E. Sievers, Znorg. Chem., 6 , 1105 (1967). (15) K. J. Eisentraut and R. E. Sievers, in “Inorganic Syntheses,” W. L. Jolly, Ed., Vol. XI, McGraw-Hill, New York, N.Y., 1968, p 94. (16) P. V. Demarco, T. K. Elzey, R. B. Lewis, and E. Wenkert, J. Amer. Chem. SOC.,92,5734(1970). (17) M. K. Archer, D. S. Fell, and R. W. Jotham, Znorg. Nucl. Chem. Lett., 7,1135 (1971).
Table I. Overall Aggregation Constants of the fod Chelates
Pn = [n-mer]/[monomerln Chloroform Carbon tetrachloride n-Hexane
EU(f0d)s log P 2 log Pa 0.2 ... 2.00 3.70
... 4.78
Pr(fod)s log 8 2 0.2
log P 3
2.53
... 6.06
3.71
...
mers in solution is believed to be due to impurities, for this peak does not appear for solutions of freshly purified compounds in purified solvents. The formation of Eu(fod), and P r ( f ~ d )aggregates ~ does not introduce additional peaks in the NMR spectra and the shift of the tert-butyl group peak is fairly constant in the concentration range investigated. This study was concerned only with aggregations in solutions of the free metal complexes. The extent of this phenomenon is very likely lowered upon the addition of an organic substrate which, in practice, is usually present in large excess. Under these conditions, the aggregation of the residual free fod chelate is low and is, thus, not necessarily reflected in the observed chemical shifts. RECEIVED for review April 3, 1972. Accepted July 19, 1972. Work supported in part by the National Science Foundation under Grant GP-30692X. Ofie of the authors (J.F.D.) also wishes to acknowledge financial support under fellowships from NATO and Patrimoine de 1’UniversitC de Likge (Belgium),
GINA-A Graphical Interactive Nuclear Magnetic Resonance Analysis Program Stephen R. Heller Heuristics Laboratory, Dicision of Computer Research and Technology, National institutes of Health, Bethesda, Md. 20014 Arthur E. Jacobson Laboratory of Chemistry, NIAMDD, National institutes o j Health, Rethesda, Md. 20014
THE AVAILABLE COMPUTER PROGRAMS (1-4) for calculating theoretical spectra are slow and cumbersome to use. Direct access to the WYLBUR text-editor system (5) on the DCRT/ Computer Center Branch IBM 360/370 batch programmed computer system at NIH enabled us to eliminate the error prone punched-card approach to the analysis program. However, in such systems the results from both the initial fit and iterative calculations are slow in forthcoming and the
Casetellano and A. A. Bothner-By, J. Chem. Phys., 41, 3863 (1964). (2) J. D’. Swalen and C. A. Reilly, ibid., 42,440 (1965). (3) R. K. Harris and C M Woodman, Mol. Phys., 10,437 (1966). (4) R. B. Johannesen, J. A. Ferretti, and R. K. Harris, J . Magn. Resonance, 3, 84 (1970). (5) S. J. Kaufmann, A. E. Jacobson, and W. F. Raub, J. Chem. Doc., 10,248 (1970); see references 2 and 3 therein.
(1) S.
Calcomp plots of the results are even slower. While the computer run normally takes 1-2 hours, the overall elapsed time including the plots normally takes several days. The usual procedure for an initial fit calculation, with a number of first guesses, would take at least a week, or, on occasion, .weeks, before a reasonable first fit is made for the experimental :spectrum. In addition to the time factor, the cost of the numerous Calcomp plots and the calculations tend to make the analysis an expensive proposition. Thus we decided to rewrite the original UEAITR program (4) for the PDP-10, a time-sharing computer. The advantages of the time-sharing computer are numerous, with the main ones being the possibility of direct and immediate interaction and a graphic display of the resulting calculation. We denote the composite of the original program, now capable of display and plotting of the initial fit and iterative calculations, as GINA (Graphical Interactive NMR Analysis).
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