1261
SPECTROSCOPY OF SULFUR IN ETHYLENEDIAMINE
Spectroscopy of Sulfur in Ethylenediamine
by Robert MacColl and Stanley Windwerl* Department of Chemistry, Adelphi University, Garden City, New York 11630 (Received September 0 , 1069)
The ultraviolet-visible spectra of green solutions of sulfur in ethylenediamine exhibit bands at 616, 400, 316, and 239 mp. The concentration and temperature dependence of these bands and their decomposition show the 616-mp band represents a different species from the others. In the lower concentration range, a shoulder appears at 266 mp, The spectra of solutions of sulfur in 1,3-diarninopropanewere studied and compared to the ethylenediamine results. The 616-mp band again showed behavior that was different from the other bands. The 616-mp band was shown to represent a negative species by moving-boundary electrophoresisexperiments. The results are interpreted in that the 616-mp band is due to sulfur radical ions (-S-Ss-S.). A scheme is presented to account for these results. Past reports of sulfur radical ions are discussed, and comparisons are made to similar systems.
Introduction Solutions of sulfur in many organic solvents contain sulfur dissolved in the form of cyclooctasulfur molecules. These solvents include benzene, alcohols, and hexane. Friedman and Kerkerlb found an almost linear relationship between the refractive indices (n) of these solvents and the extinction coefficients at a particular wavelength in the ultraviolet absorption region. The relation between extinction coefficient and refractive index was derived by Davis2 using quantum theory. He found that the extinction coefficient was 4n 4/n. Davis and actually proportional to ns NakshbendiS found this relationship was accurate for sulfur in many solvents, but amines did not obey this cubic equation. It is concluded that amines do not dissolve sulfur as cyclooctasulfur molecules. That sulfur-amine solutions contained ions was demonstrated by conductance measurements. 8,4 Hodgson, Buckler and Peters5 showed the presence of free radicals in these amine-sulfur solutions by measuring their electron spin resonance spectra. The purpose of this investigation was to study sohtions of sulfur in ethylenediamine in an attempt to better understand the nature of these solutions. The principal method used was ultraviolet-visible spectroscopy under a variety of conditions.
+ +
Experimental Section Sulfur (Baker, NP grade) was purified by the method of Bacon and FaneIli.0 Ultrapure grade sulfur from the Research Chemical Company was used without additional purification. Both types of sulfur gave identical spectroscopic results. Ethylenediamine (Baker reagent grade) and 1,3-diaminopropane (Eastman Yellow Label) were treated as follows. The amine was refluxed from reagent grade potassium hydroxide and anhydrous barium oxide under nitrogen and a pressure of about 20 Torr and then distilled over a packed 900-mm column. The
middle cut was collected. The distillation process was repeated. The flask containing the doubly distilled amine was atached to a high-vacuum system by means of a ground-glass joint. The remainder of the vacuum system was evacuated using a mechanical pump and an oil diffusion pump. Pressures obtained were usually Torr, but pressures as high as 5 X near 5 X Torr were also considered adequate. Apiezon N or T greases were used on stopcocks and joints on the highvacuum system. The amine was then allowed into the vacuum system and distilled under vacuum over lithium. The distillation proceeded as follows. A dark blue solution of lithium in amine was heated in a warm water bath, while the collection flask was placed in an acetone-Dry Ice bath. This collection flask contained additional pieces of lithiurn, After completion of the first distillation, the frozen material was immediately melted with a warm water bath and a dark blue solution formed as the lithium dissolved. Lithium was less soluble in 1,3-diaminopropane than it was in ethylenediamine. The amine was distilled into a solvent reservoir flask. Nitrogen from a nitrogen reservoir flask was used to push amine through an inlet tube into a buret as needed. The buret emptied into a reaction flask that contained a weighed amount of sulfur. The reaction flask was evacuated using the oil diffusion pump prior to the addition of the amine. (1) (a) To whom requests for reprints should be addressed; (b) H. L. Friedman and M. Kerker, J. Colloid Sci., 8,80 (1953). (2) R. E. Davis, Proc. Indiana Acad. Sci., 71, paper I1 (1962). (3) R . E. Davis and H. F. Nakshbendi, J. Amer. Chem. Soc., 84,2085 (1 962). (4) G . L. Putnam and K. A. Kobe, Trans. BEectrochem. Soc., 74, 609 (1938). (5) W. G . Hodgson, S. A. Buckler, and G . Peters, J. Amer. Chem. SOC.,8 5 , 543 (1963). (0) R . F. Bacon and R. Fanelli, Ind. Eng. Chem., 34, 1043 (1942); J. Amer. Chem. SOC.,65, 639 (1943); T.J. Murphy, W. S. Clabough and R. Gilchrist, J. Res. Natl. Bur. Std., 69A,355 (1960).
Volume 74, Number 6 March 19, 1070
1262 Two reaction flasks, both specially designed to fit into the normal cell compartment of the Cary 14 spectrophotometer, were used. The first had a 2.003mm quartz optical cell (America1 Instrument Co.) attached through a quartz-Pyrex graded seal. The second reaction flask had two quartz cells (0.105 and 0.504 mm) at about ninety degrees apart. This allowed a change in path length by simply rotating the flask in the cell compartment of the Cary 14. Neutral density filters were used when required. When the amine had been added to the sulfur, a stopcock was closed and the reaction flask was removed from the vacuum system. A glass-coated stirring bar was used to stir the solution. All the absorption bands increased for a period of time until the absorbances were constant. The solutions in this condition exhibited no spectral changes for several hours. Some showed no change for 24 hr and more. After a period of time the solutions decomposed. This stability was considerably better than that found by earlier workers, who did not use ultra high-vacuum techniques. When sulfur and ethylenediamine are initially mixed, the sulfur dissolves to give a red color in the vicinity of the solid sulfur. The red quickly changed to green. In the concentration range studied, the color of the sulfur-ethylenediamine solutions varied from light green to very dark green at 25". The temperature in the cell compartment of the Cary 14 was regulated by a Haake constant-tempera ture circulator. Most experiments were performed at 25O, and the temperature control vias =tO0.O8or better. When temperature changes were made, the absorbances of the various bands were frequently measured. When all the bands exhibited constant absorbances for a long period of time, equilibrium was assumed to have been reestablished. All the- glassware used on the high-vacuum system cleaning solution, was cleaned with an HF-"03 rinsed five times with distilled water, treated with freshly prepared aqua regia, rinsed ten times with distilled water, and dried at 140" in an oven. . The optical cells were cleaned with dilute HC1 and distilled water. The amine used in the reference cells was purified in the same manner as the amine used in the reaction flask. Apparent molar concentrations were calculated using 32.064 for the molecular weight of the various species. All absorbances reported here have been scaled, so that they represent absorbances for a l-mm cell. The moving-boundary electrophoresis experiments were based on those of Burton.' This procedure involved the careful addition of sulfur-ethylenediamine solution to a U-tube containing purified ethylenediamine. The addition was accomplished through a stopcock on the bottom of the U-tube. Thus a sharp The JOUTWL~ of Physical Chemistry
ROBERTMACCOLLAND STANLEY WINDWER
Figure 1. Visible-ultraviolet absorption spectrum of a sulfur-ethlenediamine solution (0.086 M a t 25 and 45').
boundary was established between the green solution and colorless solvent. When voltage was applied, this sharp boundary was not maintained. Either sodium chloride or potassium nitrate was present to act as an inert electrolyte, and the electrodes were platinum disks. The voltage was maintained at 150 V, and the current was less than 1mA.
Results The ultraviolet-visible spectra exhibited four absorbance bands (Figure 1). As the temperature was increased, the 616-mp band increasedin absorbance and the bands at 400, 316, and 239 mp decreased. Thus, the broad 616-mp band is probably caused by a different chemical species. The temperature dependence of the 400- and 316-mp bands was slight, about a 0.2 t o 0.4% change in absorbance for each degree, while the absorbance of the 616-mp band changed about 20% per degree. Little quantitative data was obtained on the 239-mp band, since its peak was very near the solvent cut-off wavelength.8 The concentration dependence of the electronic spectra was studied from 0.00164 to 0.151 M (Figure 2). The absorbance at 400 and 316-mp increased linearly with concentration, but the 616-mp band had a nonlinear dependence (Figure 3). At the lower concentrations, the 616-mp band had greater absorbance than the (7) E. F. Burton, Phil. Mag., 11, 425 (1906). (8) Measurements of the 239-ma band were limited to the 0.105-mm cell and concentrations between 0.0680 and 0.0870 M .
1264
ROBERT MACCOLL AND STANLEY WINDWER
03-
A A
-
0.1
-
c A 300 lmpl
Figure 5. Spectrum of the 266-mp band in a 0.00236 M sulfur-ethylenediamine solution at 25 and 45O.
616-mp and 400-mp species was not effected by this experiment. The green color migrated toward the positive electrode. Thus, it is concluded that the species responsible for both the 616- and 400-mp bands are negatively charged.
Discussion The 616-mp absorption band has characteristics which indicate that it may represent a free-radical species. Electron spin resonance spectroscopy showed an increase in the concentration of the free radicals when the temperature was raised.b The absorbance of the 616-mp band, likewise, increased when the temperature was increased. This band showed instability by decaying faster than the others. The wavelength of its absorption maximum is typical for sulfur radicals. Green, p ~ r p l e and , ~ most commonly blue (absorption maxima between 580 and 620 mp) radicals have been reported. Free radicals with unpaired spins on sulfur atoms have been postulated as the source for the blue color of sulfur dissolved in LiC1-KC1 melts,l0 sulfuroleum solutions1' and in several other paramagnetic, sulfur-containing systems.12 Thus, the 616-mp band seems to represent an anion radical. Evidence that the unpaired electrons are located on sulfur atoms comes from the esr experiments of Hodgson, Buckler, and Peters.6 Ethylenediamine was one of the amines studied. Most primary and secondary amines produced paramagnetic solutions, and tertiary amines did not. Some esr measurements made during this present study confirmed that sulfur-ethylenediaThe Journal of Physical Chemistry
I
350
CONC. x102
I
lo
Figure 6. Concentration dependence of the absorption spectrum of sulfur-1,3-diaminopropane solutions compared to sulfur-ethylenediamine solutions a t 25': A, 400 mp; B, 316 and A) mp; and C, 239 mp. The solid data symbols (0, I, represent l,&diaminopropane solutions and the open symbols (0,0, and A) represent ethylenediamine solutions.
mine solutions are paramagnetic. They found a g value of 2.030 for paramagnetic sulfur-amine solutions. Due to spin-orbit coupling, sulfur radicals exhibit characteristically higher g values (in liquid sulfur, g = 2.024) than the free electron spin g factor of 2.002.1a This high g value is considered excellent proof that, the unpaired spin is on a sulfur atom. This anion radical present in sulfur-ethylenediamine solutions is probably a sulfur radical ion
5-5,-8 The high concentrations (0.3 to 3M) needed for the esr experiments compared to those for the ultravioletvisible work indicates a low concentration of these radicals. DavisS has proposed a scheme for the initial steps in the reaction of sulfur and amines. For ethylenediamine this scheme would be (9) For example, H.E. Raford, and F. 0. Rice, J. Chem. Phys., 33, 774 (1960). (10) J. Greenberg, B. R. Sundheim, and D. M. Gruen, ibid., 29, 461 (1958). (11) For example, M.C. R. Symons, J. Chem. SOC.,2440 (1967). (12) Giggenbachle related the blue color of several sulfur-containing materiala to the presence of sulfur radicals. (13) D. M. Gardner and G. K. Fraenkel, J. Amer. Chem. SOC.,78, 3279 (1956).
1265
SPECTROSCOPY OF SULFURIN ETHYLENEDIAMINE Table I : Absorption Spectra of Cyclooctasulfur Solutions Solvent
T,OC
Inert (n-hexane, etc.)
Room temp. 25 - 35 and -40d 20 400 20
Ethylenediamineb Liquid ammonia DMF' with NazS Alkali halide melts Oleum (65%) Oleum (65%)
Xmax,
mr
260-280,225 616,400,316,239c 580 ,430,295-97 617 ,435 ,347 588,417,278 588,400,333 590 ,355
...
Color'
Canary yellow Green Green to blue-red" Blue Blue Blue Blue
Referenaes
3 This research 28,29 18 16 9
11
'Davis and Nakshbendis found sulfur-amine solutions with colors ranging from orange-yellow and red-orange to green. At 60", for solutions around 0.002 M the color is blue. For solutions around 0.002 M a shoulder emerges a t 266 mp. * At - 77", Nelson and LagowskiZBfound the peak at 580 mw had disappeared. e Solutions 0.01 M and greater are dichroic and show a blue-red color. Sohtions less concentrated are green.** D M F = dimethylformamide. 'H. Lux and E. Bohm, Chem. Ber., 98,3210 (1965).
'
NHzCHzNHz
+ Ss
NH2CH2CH2N+Hz-S.r-S-
NHzCHzCHzN+Hz-S+Y'
N+H3CH2CH2NH2
+ NHzCHzCHzNHz
(1)
FAST
+ NH2CH2CHzNHS7-S-
(2)
Equation 1 shows a nucleophilic attack on the sulfur ring causing heterolytic cleavage of the ring. Many substances react with sulfur in this manner.14 We suggest that sulfur radical ions could then arise from homolytic scission of an S-S bond. NH&H2CHzNH-S&NH2CH2CH2NH-Sz-S
9
+ *S-Ss-z-S-
(3)
Levil6 has found that amines and sulfur can react to yield N,N'-polythiobisamines. These N,N'-polythiobisamines could be formed by the dimerization of the thioamine from eq 3. I n general, this would occur as follows. 2RNH-Sz-S
*
RNHS,-S-S-S,-NHR
(4)
Since sulfur radical ions are postulated for sulfurethylenediamine solutions, the existence of these species in other systems is of great interest. The scarcity of data on these ionic radicals and the controversial nature of some of these reports makes the following brief discussion on sulfur radical ions pertinent and necessary. Giggenbachl6 found the S2- ion, having an absorbance maximum at 617 mp, in the blue solutions of alkali metal tetrasulfide in dimethylformamide. He also postulated that sulfur radical ions were present in solutions of sulfur and alkali metal sulfide in dimethylformamide, where a small amount of water was also present. I n addition to the Sz- ion, he postulated the existence of SA- and Sa- (g = 2.027). The Sz- ion shows extreme g anisotropy, and Morton reported that its resonance spectrum is undetected except a t liquid helium temperature." Lux, Benninger, and Bohrn'S studied very dilute solutions of sulfur and alkali metal sulfide in anhydrous dimethylformamide and found no sulfur ions present other than the sulfide ion.
Sulfur radical ions have been suggested for ultramarine16Jg-21and sulfur-doped alkali metal halide^."^^^^^^ Gardner and Fraenkel'g reported a g value of 2.028 for the blue ultramarine. I n selenium-doped alkali metal halide^,^^,^^ Se2- and Sea- have been observed. I n alkali metal halides doped with both selenium and Sunamoto, sulfur, SSe- was d i s ~ o v e r e d . ~Hashimoto, ~ and A ~ k found i ~ ~ an esr signal in solutions of alkali metal tetrasulfide in water that were kept under nitrogen for 30 days. They also postulated the presence of sulfur radical ions. However, they found a g value of 2.002, which is lower than the other reported values for sulfur radicals. Solutions of sulfur in liquid ammonia give results similar to those in ethylenediamine. The sulfurliquid ammonia solutions have been studied since 1873,26 and the species present are still unknown. Both liquid ammonia2' and ethylenediamine3t4dissolve sulfur to produce ionic solutions. The absorption (14) M. Schmidt in "Elemental Sulfur," B. Meyer, Ed., Interscience Publishers, New York, N. Y., 1965,Chapter 16. (15) T. G. Levi, Gam. Chim. Ital., 60, 975 (1930); ibid., 61, 286 (1931). (16) W. Giggenbaoh, J.Inorg. Nucl. Chem., 30, 3189 (1968). (17) J. R. Morton, J. Chem. Phys., 43, 3418 (1965); J. Phys. Chem., 71, 89 (1967); J. R. Morton and L. E. Vannotti, Phys. Rev., 161, 282 (1967). (18) H, Lux, S. Benninger, and E. Bohm, Chem. Ber., 101, 2486 (1968). (19) D.M. Gardner and C. K. Fraenkel, J. Amer. Chem. SOC.,77, 6399 (1966). (20) W.G. Hodgson, J. S. Brinen, and E. F. Williams, J. Chem. Phys., 47, 3719 (1967). (21) R. Bijttoher, S. Wartewig, W. Windsch, and A. Zschunke, 2.Naturforuch., 23a, 1766 (1968). (22) J. H.Sohulman and R. D. Kirk, Solid State Commun., 2 , 105 (1964). (23) J. Schneider, B. Disohler, and A. Rauber, Phys. Status SoZdi, 13, 141 (1966); J. Sulwalski and H. Seidel, ibid., 13, 169 (1966). (24) L. E.Vannotti and J . R. Morton, Phys. Lett., 24A, 520 (1967); J. Rolfe, J. Chem. Phys., 49, 4193 (1968). (25) S. Hashimoto, J. Sunamoto, and A. Aoki, Bull. Chem. SOC.Jup., 40, 2867 (1967). (26) G. Gore, Proc. Roy. SOC.,21, 140 (1873). (27) E.C.Franklin and C. A. Kraus, Am. Chem. J.,24,89 (1900). Volume 74,Number 6 March 19, 1970
1266
J. C. HINDMAN, A. SVIRMICKAS, AND M. WOOD
spectra were studied by ZippZaand Nelson and LagowskiaZ9 They found bands with maxima a t 580,430, and 295-297 mp. The temperature dependence and the manner of decomposition were similar in both solvents. Zipp studied the concentration dependence of the visible spectrum and found that at lower concentrations the 430-mp band had higher absorbances than the 580mp band, but a t higher concentrations the 580-mp band had the higher absorbances. We found the reverse relationship between the 616-mp band and the 400-mp band in ethylenediamine. Table I shows a comparison of the absorption spectra of some sulfur solutions. In both amines and liquid ammonia sulfur dissolves r e v e r ~ i b l y . ~From * ~ ~ liquid ammonia, sulfur is recovered by evaporation and from amines by dilution with aqueous hydrochloric acid. The sulfur-ammonia solutions, if they are truly analogous to the amine solutions, will be paramagnetic. Therefore, esr spec-
troscopy may prove valuable in the study of these liquid ammonia solutions. The conductance behavior of these sulfur-liquid ammonia solutions was recently reported by Zipp and ever^.^^ The concentration and temperature dependence of the visible absorption spectra of sulfur-ethylenediamine solutions indicate that the 616-mp species is involved in an equilibrium relationship. If this equilibrium was either a dimerization or a disproportionation, the sulfur radicals would form polysulfide ions. Acknowledgments. R. M . gratefully acknowledges the receipt of a NASA Predoctoral Fellowship during the tenure of this work. (28) A. P.Zipp, Ph.D. Thesis, University of Pennsylvania, 1964. (29) J. T. Nelson and J. J. Lagowski, Inorg. Chem., 7, 1292 (1967). (30) A. P.Zipp and E. C. Evers, ibid., 8, 1746 (1969).
The Spin-Lattice Relaxation of Oxygen-17 in Water1 by J. C. Hindman, A. Svirmickas, and M. Wood Chemistry Division, Argonne National Laboratory, Argonne, Illinois 60439 (Received September 3, 1969)
-
The quadrupole relaxation time, 7'1, has been measured in liquid water over the temperature range from 17 to 95'. In TI is found to be a smoothly varying function of 1/T down into the supercooled region. The temperature behavior is non-Arrhenius. The data can be fit by the empirical equation, In T I = 75.13 i 3.53 - 10.76 f 0 . 5 3 In T - 5608 + 161/T. Similar equationscan be used to express the temperature dependence of dielectric and deuteron quadrupole relaxation data in water. The question of the temperature variation of the oxygen17 and deuteron quadrupole coupling constants in liquid water is discussed. The physical interpretation of the equation used to represent the relaxation data is also discussed briefly.
Two recent experimental studies have been made of the oxygen-17 quadrupole relaxation in water.2~3GlaseP concluded, from his measurements and a comparison of the deuteron and oxygen-17 spin-lattice relaxation times with dielectric relaxation data, that there was a temperature coefficient for both the 2H and ' 7 0 quadrupole coupling constants, with a well defined minimum in the 1 7 0 coupling constant in the region of 40'. On the other hand, Garrett, Denison, and Rabideaua concluded from their data that the ' 7 0 quadrupole coupling constant was essentially independent of temperature in the interval between 5 and 95". Since in both cases the calculations of the quadrupole coupling constants were based on the assumption of a classical rotational diffusion model for the relaxation process, the data appear to be discordant. In view of the importance of the conclusions with respect to the nature of the relaxing The Journal of Physical Chemistry
species, a reinvestigation of the 170relaxation appeared in order. The present communication summarizes the results.
Experimental Section The H2170water samples were prepared from Yeda Research and Development Company water enriched to 26.94% in oxygen-17. The samples were prepared by multiple distillation on a vacuum line and sealed after gas removal by the usual freeze, pump, thaw technique. To ensure reproducibility of the results it was found necessary to boil all the glassware for several hours in (1) Based on work performed under the auspices of the U.9. Atomic Energy Commission. (2) J. A. Glasel, Proc. Nat'l. Acad. Science, 58,27 (1967). ( 3 ) B. B. Garrett, A. B. Denison, and S. W. Rabideau, J . Phys. Chem., 71,2606 (1967).
