1758
catalyzed decay of metastable species when a trace of nitric oxide was added between the discharge and the pink afterglow region. An attempt was made to deactivate excited species by placing glass wool between the discharge and the titration inlet (experiment 11). The ratio (N)No/ (N)e.S.r. decreased somewhat although not to unity. This result suggests that certain metastable species may be capable of surviving a large number of surface collisions. However, the results of this experiment were only approximate owing to the low concentration of atomic nitrogen present. A much more extensive study of this point might be of interest. In summary, the data presented indicate that the nitric oxide titration method may not be valid under all circumstances. Apparently, it does give correct measurements of nitrogen atom concentrations, provided there is a sufficiently long time interval between the discharge and the titration zone to permit deactivation of excited species interfering with the titration process. Acknowledgment. The authors wish to express their appreciation and gratitude to Professor C. A. Winkler for initiating this work while visiting this laboratory and for his continued interest and assistance. We also wish to thank Drs. C. Barth, A. Hildebrandt, and A. Westenberg for valuable discussions. This work was sponsored by the National Aeronautics and Space Administration under Contract NAS 7-100.
NOTES
In 1,2-dinitrotetrafluoroethane (I) the strong inductive effect of the fluorines should withdraw electrons from the nitro groups, making them poorer electron donors and better electron acceptors and thus permitting stronger interaction with electron-donor solvents. Such solutions have now been examined by means of absorption spectra. A 0.05 M solution of methanol in 1,2-dinitrotetrafluoroethane gives a single hydroxyl band a t 3650 cm.-l (e 75, Avl” 26 cm.-l) in a high-resolution spectrophotometer, indicating that there is no detectable hydrogen bonding under these conditions. It is interesting that there is actually a blue shift of the hydroxyl band in this solvent compared to carbon tetrachloride solutions of methanol, which have vmsx 3640 cm.-l. In solutions of anhydrous alcohols and amines 1,2dinitrotetrafluoroethane (I) is sufficiently stable so that absorption spectra can be determined. The spectra remain unchanged over a period of 1-2 hr., and characteristic stretching bands of the nitro group appear in the expected region (Table I). I n contrast with nitromethane, the asymmetric NOz band is shifted slightly by solvents, and it is probably significant that the shift from carbon tetrachloride to methanol is toward higher frequencies. The half-band widths are largely unaffected (Av’” = 20 f 2 cm.-I). The relative absorption intensities, on the other hand, are definitely greater in electron-donor solvents such as triethylamine although the relationship does not appear to be simple.
1,2-Dinitrotetrafluoroethane, Absorption
Spectra and Bonding with Solvents1
by H. E. Ungnade, E. D. Loughran, and L. W. Kissinger Universitg of California, Lo8 Alamos Scientific Laboratory, Los Alamos, New Mexico (Received November 22, 1964)
Previous investigations of the ultraviolet absorption spectra of aliphatic nitro compounds have established that abnormally high absorption intensities in the region of 280 mp occur in alcohol, ether, and amine sol~tions.~-~ Since the effects are greatest with strong electrondonor solvents, they have been attributed to an interaction in which the nitro compounds act as electron acceptors. I n the case of alcohols the interaction is further complicated by the existence of weak hydrogen bonds between hydroxyl and nitro group^.^ The Journal of Physical Chemistry
Table I : Infrared Absorption Bands of 1,2-Dinitrotetrafluoroethane Solvent
Gas
CCL Dioxane MeOH CHC1, CH2C12 EtgN
v,.-Noz
1623 1610 1610 1615 1609 1610 1611
. .. 1100 1200 1370 1260 1440 1790
~ 2 ’ 2
v.ym-~Op
21 18 20 22 20 18 19
1276 1271 1270 1285 1274 1280
...
B
~v’/zb
... (220)
36 (40)
(350) (290) (210)
(30) (30) (30)
..,
...
a Fused band with considerable overlap in solutions. values in parentheses are only estimates.
...
...
The
(1) This work was performed under the auspices of the U.S. Atomic Energy Commission. (2) N. S. Bayliss and C. J. Brackenridge, J . Am. Chem. SOC.,77, 3959 (1955). (3) H.E.Ungnade, E. D. Loughran, and L. W. Kissinger, J . Phys. Chem., 64, 1410 (1960). (4) H.E. Ungnade, E. M. Roberts, and L. W. Kissinger, ibid., 68, 3225 (1964),and cited references.
NOTES
1759
Ultraviolet absorption spectra have been determined for I in the gas phase and in various solvents. The results (Table 11) show that the intensity of the n + T* band near 283 mp is somewhat greater than in simple dinitro paraffins and the maxima occur at longer wave lengths.a A similar behavior has been observed with other fluoroalkyl nitro compound^.^ The band is shifted to the blue in going from hexane to alcohol
I
GAS
Table 11: Ultraviolet Absorption Bands of 1,2-Dinitrotet,rafluoroethane Solvent
hn,
Gas Hexane Cyclohexane' CC4 MeOH
282.5 286 286 286 283 283 250se 254" 275 290 297
HzO Et20 Dioxane HexNHz PrzNH EtaN
e
85.1 111 118 119 137
. . .d
310 640 485 447 388
',,A,
c
...
- cob
...
...
...
...
...
...
... ... ... 255 270 300 330
Figure 1. Ultraviolet absorption spectra of 1,2-dinitrotetrafluoroethane as gas and in several solvents.
...
... ...
*..
LOG
CF2N0,CF2N0,*n
1
El
N =:;M \IN L / HEXANE
E-€()
2.0
600 405 343 348
Corrected for the absorption of both components. eo is the molar absorptivity of 1,2-dinitrotetrafluoroethane in hexane a t Amax,. The corresponding values for 2,3-dimethyl-2,3dinitrobutane are 284 mfi ( E 75.7): H. E. Ungnade and L. W. The compound is Kissinger, J . Org. Chem., 22, 1088 (1957). very insoluble in water; a saturated solution a t 25' was used to Overlapping bands (s sign8es shoulder). determine A=,.
and water solution but there is no longer a blue shift compared to gas. There is a pronounced increase in the integrated intensities in the order R H < ROH < ROR < RNHa < R2NH < RsN. Since the infrared absorption spectra do not show any hydrogen bonding with alcohols, the intensity increase in this case is ascribed to the electrondonor properties of these compounds. Ethers and amines being stronger electron donors give still larger intensity increases and simultaneous shifts of t4e absorption band to longer wave lengths. The shifts are so large in this case that the triethylamine solution absorbs in the visible region, and the solution is bright yellow. The absorption spectrum at longer wave lengths actually consists of two fused bands, the still existing nitro n --+ T* bands and new bands which bear some resemblance to charge-transfer bands owing to their intensity and extraordinary broadness (Figure 1). The interaction between the dinitro compound (I) and triethylamine (11) has been investigated further. A dilute solution of 1 and two molar equivalents of 11 in hexane shows virtually no interaction. The uht+
1
1.0
mox I 300
330
I 350
Figure 2. Ultraviolet absorption spectra of approximately 0.03 M l,>dinitrotetrafluoroethane and various concentrations of triethylamine in hexane.
violet absorption spectrum above 250 mp is essentially that of I. With increasing concentration of I1 a new broad band appears at longer wave lengths. This band overlaps with the nitro band at 286 mp up to wave lengths of about 370 mp. The spectrum of the perturbation can be obtained by subtracting the absorptivity of the nitro band at each point below 370 mp. This gives a series of broad curves with A,, 330 mp (Figure 2).6 For any given wave length in the curves in Figure 2 the absorbance in a 1-cm. cell is found to be strictly proportional to the product of the molar concentrations of the components. A Benesi-Hildebrand plot gives a straight line passing through the origin, from which K = 0 and E = ~ 0 . It is concluded that the interaction is one of contact charge transfer, as is the case for nitromethane-triethylamine. a NI4 n.m.r. determinations (by M. Alei and L. 0. (5) R. N. Hazeldine, J. C h m . SOC.,2525 (1953). (6) The amine absorption was for by using Ltnamine solution of identical concentration in the reference beam.
Volume 69,Number 6 May 1966
NOTES
1760
Morgan) show single bands for I and I1 at +1.46 and +5.75 gauss relative to aqueous nitrite = 0. These bands remain unchanged in a 1:2 molar mixture of I and I1 as far as chemical shift and band width are concerned. The results are consistent with the type of interaction postulated and eliminate other possibilities such as unpairing of electrons or formation of ions.
Experimental Materials. Spectro Grade solvents were used in all determinations. The l12-dinitrotetrafluoroethane(I) was 98% pure as determined by mass spectrographic analysis.' Determination of Absorption Spectra. In view of the volatility of I, the liquid was weighed in stoppered volumetric flasks, immediately cooled to Oo, and mixed with the appropriate solvents which were precooled. The solutions were warmed to 20°, made up to the correct volume, and run within a few minutes. Infrared absorption spectra were determined with a Perkin-Elmer Model 421 spectrophotometer and ultraviolet absorption spectra with a Cary Model 14 instrument. Solutions of I in moist triethylamine changed with time. An isosbestic point was observed at 305 mp, and the solution contained nitrite ion indicating that a nitro group was eliminated. In anhydrous solutions of amines the ultraviolet absorption spectra showed no appreciable changes during 2 hr., and there was no nitrite ion in solution. The compound could be recovered unchanged in 80% yield from its triethylamine solution by pouring on ice and hydrochloric acid and extracting with carbon tetrachloride. Because of the toxicity of 1,2-dinitrotetrafluoroethane, it must be handled with caution.
and has been used to determine the rotational conformation of the CHO group in a number of substituted benzaldehydes.lt2 Long-range aldehyde couplings have also been observed in furan and thiophenealdehydes,* but the possible stereospecilicity of these couplings has, as yet, not been established. In order to investigate this we have studied the p.m.r. spectra of 2furanaldehyde at temperatures down to -100". At room temperature the aldehyde signal is split owing to spin coupling with the ring hydrogens in the 4- and 5positions (in dimethyl ether solution at +35" J C H O - 4 = 0.31 and J c ~ 0 - 5 = 0.66 c.P.s.). Both these spin couplings have recently been shown to carry the same sign as the ring coupling constant, J 4 , 6 . 4 At about -60" the aldehyde signal, as well as the signal from the ring hydrogen in the 3-position1becomes broadened, and at -80" two signals from the aldehyde group and two signals from the ring proton in 3-position-both pairs in the intensity ratio 0.13 :0.87-are seen. The smaller aldehyde signal is found at 0.17 p.p.m. toward lower fields than the larger signal, and the smaller signal from the 3-hydrogen is found at 0.15 p.p.m. toward higher fields than the larger signal. From considerations of the diamagnetic anisotropy of the >C=O group, which causes deshielding of protons in the trigonal plane16 it seems plausible that the stronger signals may be assigned to the CHO proton and H(3-) proton in the planar rotational isomer I, and the corresponding set of smaller signals may be assigned to isomer 11.
1
1
H
I (7) The authors are indebted to Dr. G. Dorough of the Livermore Radiation Laboratory for a generous sample of this compound. 1
II
0
I1 Rotational Isomerism and Stereospecific Long-Range Spin Coupling in 2-Furanaldehyde
by Kjell-Ivar Dahlqvist and Sture Forsen Research Group for Nuclear Magnetic Resonance, Division of Physical Chemistry, Royal Institute of Technology, Stockholm 70, Sweden (Recefved October 18, 1964)
The stereospecificity of the long-range aldehyde spin coupling observed in proton magnetic resonance (p.m.r.) spectra of benzaldehydes is well established The Journal of Physical Chemistry
The predominance of the isomer I may possibly be rationalized in terms of the mutual repulsion of the two oxygen atoms. (1) 8.Forsen and B. Akermark, Acta Chem. S c a d . , 17, 1712 (1963). (2) G. J. Karabatsos and F. M. Vane, J . A m . Chem. SOC.,85, 3886 (1963). (3) For a recent extensive review of long-range proton spin couplings see S. Sternhell, Rea. Pure Appl. Chem., 14, 15 (1964). (4) R. A. Hoffman, B. Gestblom, S. Gronowitz, and S. Forsen, J . Mol. Spectry., 11, 454 (1963). (5) L. M. Jackman, "Applications of Nuclear Magnetic Resonance Spectroscopy in Organic Chemistry," Pergamon Press, Ltd., Oxford, 1959, Chapter 7.
