4708
NOTES
A measurement of the a form has also been carried out at 25" to detect any temperature dependence. We measured the height of the nuclear quadrupole resonance lines for the a and p form before and after the y radiation with a dose of 1.1 X lo8 R. I n order to avoid instrumental variations the height of the line of the sample under study was always compared with that of a standard sample measured afterward.
Results and Discussion We found that the height of the nuclear quadrupole resonance line corresponding to the a form decreased by about 21% and that of the p form by about 10% for the same dose of irradiation. This difference in the radio sensitivity cannot, probably, be due to a temperature effect (since the measurement of the a form at 25" showed no difference from that at 20") ;hence it can only be attributed to a difference in the crystal structures of the two forms. In the triclinic' arrangement all the molecules are parallel; the monoclinic arrangement is achieved by pivoting the molecules in every second layer along a about their respective centers in the ab plane followed by a translation of b / 2 . The intermolecular contacts of C1. C1 is 3.46 A for the @ form and 3.85 A for the a form.8 The resulting structure of the a phase is more open, and, consequently, the chlorine atoms freed from the rupture of the C-C1 bond can migrate more easily and in larger quantities into the space between the molecular planes than in the p form. Another reason for the greater stability of the @ form lies in the fact that in this structure the C1. .C1 intermolecular contacts are shorter than in the a form. These results confirm the fact that purely crystallographic factors can affect radioresistance in the solid state. Work is in progress at liquid nitrogen temperature where a third phase has been reported9 and in other polymorphic systems also. Acknowledgment. We wish to thank Mr. M. Voudouris for his valuable contribution in the development of the instrumentation and for technical assistance.
-
(7) J. Honsty and J. Clastre, Acta Crystallogr., 10, 695 (1957). (8) U. Croatto, S. Bezzi, and E. Bua, ibid., 5,825 (1952). (9) D. E. Woessner and H. S. Gutowsky, J . Chem. Phys., 39, 440 (1963).
The Effect of Neighboring Magnetic Anisotropy on N H Proton Chemical Shifts by B. M. Fung Contribution No. 372 from the Department of Chemistry, Tufts University, Medford, Massachusetts 02166 (Received August 2, 1968)
The effect of the magnetic anisotropy of a neighboring group on the proton chemical shift of acetylene, hydroThe Journal of Physical Chemistry
gen halides, several binary hydrogen compounds, and hydrocarbons has been studied.1-3 We wish to discuss briefly the application of the theory of neighboring magnetic anisotropy on the NH proton chemical shift on several compounds upon protonation and coordination to Co(II1). For dilute solutions of ammonia and amines in dimethyl sulfoxide (DMSO), the NH proton chemical shift did not change with concentration. This indicated that the self-association between the solute molecules was not appreciable. Therefore the data can be treated by considering only 1: 1amine-solvent complexes. For dilute solutions of diamagnetic salts in the same solvent, the effect of ion pairing on chemical shift is not very important4and will be neglected. The replacement of a hydrogen by a methyl group causes the NH proton signal to move downfield. The change in the series of ammonia, methylamine, and dimethylamine was fairly regular (Table I). For ethylenediamine, which has two functional groups, the NH proton signal lay between those of the primary and the secondary amines. Pyrrolidine is a secondary cyclic amine. Its NH proton signal appeared a t considerably lower field than a similar open-chain secondary amine, dimethylamine. The shift is larger than that in open-chain and cyclic alkanes. In the alkanes, the anisotropy of the C-C bond in the p position is responsible for the difference in shielding.8 It is possible that in amines the anisotropy of the C-C bond is enhanced by the neighboring nitrogen atom, which has a lone pair of electrons. Other structural differences between the two kinds of amine may also contribute to the difference in chemical shift. Upon protonation, the NH proton resonance experienced a large downfield shift of about 6.5 ppm for ammonia and the open-chain amines. For pyrrolidine, the change was considerably smaller (Table I). The 14Ncoupling (J" = 52.0 =t0.5 Hz) showed up only for the most symmetrical ammonium ion; the "proton in all other ions appeared as a slightly broadened singlet. The downfield shift of the NH proton is probably caused by three main factors: decrease in electron density, reduction of magnetic anisotropy due to increased symmetry, and enhanced hydrogen bonding with the solvent. For ammonia, if we consider a 0.25 positive charge being added to each hydrogen in forming the ammonium ion, the contribution to the local diamagnetic shielding would be -5.35 X 10-6.5 (1) J. A. Pople, PTOC.Roy. Soo., A239, 560 (1951). (2) H.M. McConnell, J . Chem. Phys., 27, 226 (1957). (3) A. A. Bothner-By and C. Naar-Colin, Ann, N . Y . Acad. Sci., 70, 833 (1958). (4) J. C. Fanning and R. 8. Drago, J . Amer. Chem. SOC.,90, 3987 (1968). (5) J. A.Pople, W. G. Schneider, and H. J. Bernstein, "High-Resolution Nuclear Magnetic Resonance," McGraw-Hill Book Co., Inc., New York, N. Y.,1959,p 175.
NOTES
4709
Table I: The NH Proton Chemical Shift of Ammonia and Some Amines Dissolved in Dimethyl Sulfoxide A6 (from free base), ppm
c
S (from TMS), ppm
Ammonia Methylamine Ethylenediamine Dimethylsmine Pyrro1idin.e
-0.63 f 0 . 0 2 -1.09 h 0 . 0 2 -1.26 5C0.02 -1.45 f 0 . 0 2 -2.28It0.04
Protonationa
-6.45 -6.34 -6.52 -6.67 -6.07
f0.02 f 0.02 5C 0.02" f0.02
Coordination t o Co(II1)
-2.80 f 0.02b -4.30 f 0.02d
5C0.04
a Perchlorate salts. Hexaamminecobalt(II1). ' Diprotonated. Tris(ethylenediamine)cobalt(III), averaged for axial and equatorial protons: B. M. Fung, J. Amer. Chem. SOC.,89, 5788 (1967).
The removal of the neighbor anisotropy changes the These two factors acshielding by -0.25 X count for the bulk part of the observed downfield shift. For the effect of hydrogen bonding, the ammonium ion is expected to be a better proton donor than ammonia; however, the contribution to the deshielding of proton is difficult to estimate. A downfield shift of the order of 1 ppm does not seem too unreasonable in the highly basic solvent, DMSO. The argument can be extended to explain the protonation of the open-chain amines. The NH proton of p,yrrolidine has a smaller downfield shift upon protonation. As a result, the over-all h"proton chemical shifts (from tetramethylsilane (TMS)) for dimethylammonium ion and pyrrolidinium ion (ca. 0.2 ppm) do not differ very much. The bonding of the lone electron pair to a proton apparently makes the systems behave more like normal alkanes. When coordinated to Co(III), the protons of ammonia experienced a downfield shift of -2.80 ppm (Table I). The main cause for this is again the decrease of electron density around the hydrogen atom. Cotton and Haas estimated that about 87% of the positive charges in the complex are distributed to the ligands.6 From this estimation it may be considered that a positive charge of 0.14 is added to each hydrogen in forming the hexaammine cobalt(II1) ion; then the proton shieldThe contribution of ing is decreased by 3.0 X magnetic anisotxopy is probably not very large: although Co(II1) has low-lying excited states, it is octahedrally symmetrical in the hexaammine complex. There may also be a small effect due to the increase of hydrogen bonding between the NH proton and the solvent. Upon coordination to Co(III), the N H protons in ethylenediamine move to a considerably lower field than those in ammonia. The charge distribution in the two complexes is expected to be very much the same. The effect of neighbor anisotropy, however, would be quite different, which may be the cause of further deshielding for ethylenediamine. First, the symmetry of the complex is lowered from O h to DB in the trisene complex. The low-lying excited states of the complex may then have significant contribution to the magnetic anisotropy of the cobalt-nitrogen bond. Second, ethylene-
diamine is a bidentate ligand which forms a five-membered ring with the cobalt atom. The Go-N bond for ethylenediamine cannot rotate freely as it does for ammonia. The effect on the average anisotropy seen by the hydrogen is similar to the case of C-C bonds in aliphatic compounds. Estimated from the data in Table I, AXC~-Nis probably one order of magnitude greater than AXC-C. This may also be the cause of the larger chemical shift (ca. 0.67 ppm) between the axial and equatorial N H protons' in the chelate than the shift between the CH protons (ca. 0.40 ppm) in cyclohexane. The difference in the NH proton chemical shift of ammonia and ethylenediamine was also observed which contains for cis- and trans-[C~(en)~(NH~)~]~+,* both kinds of ligand in the same molecule. This fact indicates that the effect of ring formation is the more import ant one. Jolly, et al., found that, in a number of cobalt(II1)pentaammine complexes, the ammonia molecule trans to the sixth ligand (X) has a higher field proton chemical shift than the ammonia molecules cis to X.9 They Bhowed that the effect is additive by studying the proton resonance of several cobalt(II1)-tetraammine complexes. This effect can be readily explained by the difference in the magnetic anisotropy of the Co-X bond as seen by the protons in trans and cis positions. The length of the vector connecting the protons and the center of the Go-X bond is larger for the trans position than the cis position. However, the angle (e) between this vector and the Go-X bond of usual bond distance is such that, during the rotation of the ammonia molecules about the Co-N bond, (3 cos2 e - 1) stays constant for the trans-protons but changes sign for the cisprotons. Therefore, on the average the contribution of the magnetic anisotropy to the shielding of both kinds of proton are of the same order. Interesting enough, for X = NOz-, there is no difference in the chemical shift for the two kinds of protons. This is (6) F. A. Cotton and T. E. Haas, Inorg. Chem., 3 , 1004 (1964). (7) B. M.Fung, J . Am. Chem. Soc., 89, 5788 (1967). (8) 8. T.Spees, Jr., L. J. Durham, and A. M. Sargeson, ibid., 5, 2103 (1966). (9) W.L. Jolly, A. D. Harris, and T. S. Briggs, ibid., 4, 1064 (1965). Volume 7.9, Number I S
December 1968
4710 probably because the Co-X bond in that case is not too different from the other Co-N bonds, as far as the hydrogen atoms are concerned.
Experimental Section All chemicals were analytical grade. Dimethyl sulfoxide was dried over calcium hydride for 24 hr or more and was vacuum distilled over Molecular Sieve 3A. Ammonia, methylamine, and dimethylamine were passed through a 9-ft column packed with iVolecular Sieve 3A into DMSO to make up solutions. The concentration was determined by diluting the DMSO solution with water and titrating with standard HCl solution. Pyrrolidine and ethylenediamine were directly weighed for making up solutions in DMSO. The concentrations of the amines were 0.1-1.0 N ; in this range the NH proton chemical shift of each amine did not change with the concentration within experimental error. The perchlorate salts were prepared by passing the corresponding gases (ammonia, methylamine, and dimethylamine) or adding the liquid amines (pyrrolidine and ethylenediamine) to solutions of perchloric acid in ethanol, filtering, washing with ether, and drying under vacuum. The cobalt(II1) complexes were prepared according to standard procedurelo and were recrystallized from water. The concentrations of the salt solutions were 0.2-0.5 N . The proton nmr spectra of these solutions did not change with concentration. Proton nmr spectra were taken with a Varian A-60A spectrometer a t 35". Acknowledgment. This work is supported by the National Institutes of Health. (10) W. C. Fernelius, Inorg. &n., 2, 221 (1946).
Conductance of Thallous Nitrate in Dioxane-Water Mixtures at 25" by Alessandro D'Aprano' and Raymond M. Fuoss Sterling Chemistry Laboratory, Yale University, New Haven, Connecticut 06680 (Received August 6 , 1068)
Ion association in aqueous solutions of the alkali halides is undetectable by conductance a t concentrations less than about 0.04 N , but if the range of concentration is extended to 0.1 N (the upper limit of applicability of the present theory), association constants of the order of unity are found.2 Nitrates have been assumed to be more highly associated than halides. I n order to test this assumption, we have measured the conductance of a series of nitrates in dioxane-water The Journal of Physical Chemistry
NOTES mixtures a t 25". Here we present the results for thallous nitrate. I n water, the association constant was found to be 3.2 f 0.1, with a limiting conductance of 146.20. From the latter, the single-ion conductance of the thallous ion is 74.75, in excellent agreement with an earlier value3 of 74.71.
Experimental Section Fischer's Purified grade of thallous nitrate was dried for 24 hr at 100"; portions were then weighed in platinum boats on the microbalance to make the initial solutions for the conductance determinations. Dilutions were made by weight and were calculated to volume concentrations c (equiv/l.) = po(l yw), where po is the solvent density and w is the weight concentration of the salt (equiv/kg of solution). For thallous nitrate in water, y = 0.232, from a density of 1.035 g/ml a t w = 0.1638. For solvent mixtures 2,3, and 4, y = 0.212, 0.190, and 0.141, respectively. For mixtures 5-8, where the concentrations were less than 0.007, we used y = 0.14. Water was laboratory supply distilled water, boiled vigorously and then cooled under nitrogen; its conductance was (1-2) X 10-6 mho. Dioxane was refluxed under nitrogen over potassium hydroxide at least 24 hr and was then distilled under nitrogen from silver nitrate just before use. (It had been found that dioxane distilled from potassium hydroxide still contained a trace impurity, undetectable by vaporphase chromatography, which reduces silver and thallous nitrate solutions to give colloidal or mirror metal.) Viscosities of the solvents were determined in an Ubbelohde viscometer whose water flow time was 473.1 sec. Dielectric constants were determined at 1 Four conductance cells were used, with constants 5.1313 h 0.0001, 0.85696 f 0.00002, 0.51361 f 0.00004, and 0.14464 f 0.00001. The first three were calibrated using potassium chloride solutions.6 The fourth was calibrated by comparison with the second and third, using 0.0024 N tributylammonium pictrate in 2propanol. The conductance bridge and general technique were as described by Lind and F u o ~ s . Solvent ~ properties are summarized in Table I, where the p is the density, D is the dielectric constant, 1007 is the viscosity in centipoises, and uo is the solvent conductance. The conductance data are summarized in Table 11, where c (equiv/l.) is concentration and A (cm-2 ohm-l equiv-') is equivalent conductance.
+
(1) On leave of absence from the University of Palermo, Palermo, Italy. (2) R. M. Fuoss and K. L. Hsia, Proc. Nat. Acad. Sei. U.S., 57, 1550; 58, 1818 (1967). (3) R. A. Robinson and C. W. Davies, J . Chem. Soc., 139,574(1937). (4) J. E.Lind, Jr., and R. M. Fuoss, J . Phys. Chem., 65, 999 (1961). (5) J. E.Lind, Jr., J. J. Zwolenik, and R. M. Fuoss, J . Amer. Chew. Soc., 81, 1657 (1959).
