HYDROGEN BONDIXG IIN NITROCOMPOUNDS
2.162 g. (90%) of product melting a t 113.7-114.7'. Anal. Calcd. for (&H57045S: N, 2.92; S, 6.68. Found: ?;, 2.92; S, 6.'74. Calculalion of Activity of B. Activity of B was ca1culated from eq. 14 and 16 by use of an IBhl 1620 computer.
3225
Acknowledgment. This investigation was supported in part by Undergraduate Science Education Program Grants GE-141 and GE-2920, National Science Foundation. The authors gratefully acknowledge the assistance of Nessrs. Joel Blatt and David Solan in the preliminary stages of this investigation.
Hydrogen Bonding in Nitro Compounds'
by H. E. Ungnade, E. M. Roberts, and L. W. Kissinger University o f California, Lo8 Alamos Scientific Laboratory, Los Alamos, N e w Mexico (Received J u n e 18, 1964)
~
~~~~~~
Experimental evidence from infrared and n.m.r. spectra indicates the existence of a weak hydrogen bond between nitro groups and alcoholic hydroxyl groups. Other findings from ulbraviolet absorption spectra can be accounted for by an interaction between hydroxyl oxygens and nitro groups.
An investigation of the absorption spectra of 8nitro alcohols has shown that the absorption bands due to the nitro group cam be explained without assuming an internal hydrogen bond in these compounds,2 and preliminary examinations of the infrared spectra of nitromethane solutions of alcohols have given evidence for an interaction but no satisfactory explanation of the type of bond.2 The small shift of the alcohol hydroxyl band from carbon tetrachloride to nitro-. methane solution has been regarded as doubtful evidence for a hydrogen bond in the past.s More recent investigators disagree on its significance. Thus, the hydroxyl band of water in nitromethane is said to fall on the Kirkwood-Bauer-Magat plot and to deviate from such a plot.6 Careful studies of frequency shifts and band widths for Eieveral A-H stretching modes in a number of solvents have established, however, that there is a smooth ]progression of these parameters between polar and nonpolar solvents and, therefore, an interaction which differs from one solvent to another in degree rather than kind.6 The problem has been reinvestigated by different methods. Since absorption intensities and band widths
are generally more sensitive to hydrogen-bonding effects than frequency shifts, these have been determined for the hydroxyl bands of four alcohols in several solvents under comparable conditions. The results (Table I) show that, while the shift of the methanol hydroxyl band from carbon tetrachloride to nitromethane is small, the apparent molar absorptivity and half-band width are increased. Similar increases are observed with nitrobutane and nitrobenzene and a much larger one with triethylamine. Better proton donors such as heptafluorobutanol give more intense absorption (e X Av''') while poorer proton donors like t-butyl alcohol have weaker hydroxyl bands in (1) This work was performed under the auspices of the U. S.Atomic Energy Commission. (2) H. E. Ungnade and L. W. Kissinger, Tetrahedron, Suppl. 1 , 19, 121 (1963). (3) W. Gordy, J . Chem. P h y s . , 7, 93 (1939).
(4) E. Greinacher, W. Luttke, and R. Mecke, 2. Elektrochem., 59, 23 (1955). ( 5 ) P. Saumagne and M. L. Josien, Bull. 8oc. chim. France, 813 (1 958). (6) L. J. Bellamy, H. E. Hallam, and R. L. Williams, T r a n s . Faraday Soc., 54, 112 (1958).
Volume 68, X u m b e r I 1
November, 1964
3226
H. E. UNGNADE, E. M. ROBERTS, ASD L. W. KISSINGER
nitromethane. A plot of band shifts of the methanol hydroxyl bands against half-band widths for nitro compounds and other solvents (including established proton acceptors) is a straight line. The combined evidence indicates that nitro compounds form weak hydrogen bonds with alcohols.
CM-' 3 6 4 0 3610
Alcohol
MeOH MeOH MeOH MeOH MeOH n-BuOH n-BuOH t-BuOH t-BuOM n-CaF7CHzOH n-CaF7CHzOH
CCI, MeKOz BuXOz PhNOz EtaN CCln MeXOz CCla MeN02 CCla MeNOz
O.IM METHANOL IN CCl,
a
0.6
Mole/l.
ThickneB8, Y, om. cm.?
0 0 0 0 0 0 0 0 0 0 0
1.00 0.05 0.05 0.05 0.05 1.00 0.05 1.00 0.05 1.00 0.05
005 1 1 1 1 005 1 005 1 005 1
3640 3600 3600 3600 3240 3636 3595 3616 3578 3616 3544
Avl/Z, e
A
cm.-1 om.-'
0.4
0 66 25 110 60 40 100 60 40 106 60 40 120 170 400 68 26 0 100 60 41 0 72 19 80 52 38 0 130 25 104 116 72
Band shift.
A study of the methanol hydroxyl band in carbon tetrachloride with varying amounts of nitromethane with a high-resolution infrared spectrophotometer has established that the shift of the hydroxyl band from carbon tetrachloride to nitromethane is actually due to the formation of a new band (Fig. l).' The infrared absorption bands in 0.1 M methanol solutions can be assigned to monomer, tetramer, and polymer, and it is observed (Fig. 1) that the addition of nitromethane causes the polymer band to be decreased in intensity while the tetramer band remains largely unchanged. I n more concentrated solutions (0.5 M) so much polymer is present that the small tetramer band can no longer be observed. N.m.r. studies of these same solutions show that the nitromethane peak remains unchanged relative to the methyl peak of the methanol and that, therefore, in neither molecule are the methyl protons involved in any interaction. I n 0.5 M methanol solutions, the hydroxyl proton is shifted upfield by about 50 cycles as the concentration of nitromethane is increased from 0 to 3 44. I n 0.1 M methanol, the hydroxyl proton is shifted slightly upfield upon the first addition of nitromethane. For concentrations greater than 0.1 M the shift is downfield (Fig. 2). The upfield shifts of the hydroxyl protons are in agreement with the assuniption that hydrogen-bonded methanol polymer is depolymerized to form the complex between methanol and nitromethane. There is The Journal of Physical Chemistry
3350
CONTAINING 0 - 3 . O M McNO,
0.B
Table I : Hydroxyl Bands of Alcohols Solvent
3440
0.2
3 700
3300
3500
CM-'
Figure 1. Hydroxyl stretching bands in methanol solutions.
I
I
I
O.SM MsOtl
-
/
O
1
. O /
+/
I
I
1.0
I
I
I
20 MOLESIL. MsN02
Figure 2. Chemical shifts of the hydroxyl proton in methanol solutions with reference to the methyl peak of methanol.
such an excess of polymer in the 0.5 M solution that this trend is not reversed. I n the 0.1 114 solutions, the downfield shift' is due to the removal of monomer (7) Since the completion of these investigations, high-resolution spectra have been determined also by P. J. Krueger and H. D. Mettee [Can. J . Chem., 42, 288 (1964)] and by W. F. Baitinger, P. von R. Schleyer, T. 8. 9. R. Murty, and L. Robinson (the authors are indebted t o Dr. P. von R. Schleyer for a preprint of their manuscript).
HYDROGEN BOXDING m NITROCOMPOUNDS
3227
(for complexing with the nitromethane) while the concentration of the tetramer remains approximately constant. The addition of nitromethane to methanol solutions has the additional effect of considerably narrowing the initially broad methyl and hydroxyl resonance peaks. This line-narrowing is regarded as evidence that the nilxomethane-methanol complex is favored over the methanol polymer. The infrared data further indicate that the complex is intermediate in strength between methanol tetramer and polymer. Other experimenta,l findings confirm that the interaction is weak. The frequency of the as-nitro band in nitroinethane is virtually identical in methanol and carbon tetrachloride (Table 11). The apparent molar absorptivities are affected by solvents but the change is in the same direction for a tertiary amine which cannot form ab hydrogen bond with the nitro group and an alcohol which can. It is, therefore, unlikely that studies of the nitro stretching frequencies will give information about hydrogen bonding in nitro compounds.
Table 11: Nitro Bands of Nitromethane V
Y
Solvent
as-KO2
CCla CHCla CHzClz EtaN
1563 1563 1563 1563 1562
MeOH
B
550 600 610 412 425"
sum-NO2
e
1373 1374 1374
110 140 140
This value m a y be somewhat low because of solvent absorption near t h e maximum. It was determined in 0.002-cm. cells.
Further information about the interaction between hydroxyl and nitro groups has been derived from a study of the low intensity absorption band of nitro compounds near 280 mp. This band is quite sensitive to solvents, with respect to both position and intensities. New measurements of this band in water, methanol, and hexane (Fig. 3) confirm the blue shift with increasing solvent polarity with reference to hexane and gas.a I n the casle of pyridazine and benzophenone the blue shift of the n 4 T* transition has been traced to hydrogen bonding with the solvent. Nitromethane does absorb a t the shortest wave lengths in the solvents which can forrn the strongest hydrogen bonds, L e . , water and heptafluorobutanol (Table 111). There is, however, a simultaneous decrease in the electron density on the oxygen of these solvents, which also may affect the position of the bandslo For this reason the
1
I
NITROMETHhNE IN
250
Figure 3.
I. HpO
2 6 9 (15.3)
2. MoOW
2 7 2 (16.91
3. CgH,4
278
300
(IS.9)
mP
Ultraviolet absorption spectra of nitromethane.
absorption intensities of the ultraviolet band have been examined. It was established previously that dioxane can increase the intensity of this transition over "inactive" solvents such as hydrocarbons" and that nitrogen-containing solvents give even larger increases which fall in the order of their electron-donor tendencies.12 The interaction which occurs can be recognized by an increase in the intensity of the 280-mp band, a small apparent red shift of the intense band at shorter wave lengths,13 and a pronounced increase in the absorption at the minimum (near 250 mp). We have now observed that an increase in the electron-donor properties of ethers similarly enhances the complex formation with nitro compounds (Table 111), thus, the last two criteria for such an interaction give the ether series (C1CH2CH2)20< EtnO < (MeOCH2CH2)20< dioxane.I4 (8) N. 8. Bayliss and E. G. McRae, J . Phgs. Chem., 58, 1006 (1954). (9) G. J. Brealy and M. Kasha, J . Am. Chem. SOC.,77, 4462 (1955). (10) Such a change could be expected if the solvent were to alter the electron density on the nitrogen. It has been shown that frequency shifts in RNOZare in the Baker-Nathan order: Me > E t > GPr > t-Bu [A. Balasubramanian and C. X. R. Itao, Chem. Ind. (London), 1025 (1960)l. A similar order was found in the present investigation, both in polar and nonpolar solvents (Table 111).
(11) N. S. Bayliss and C. J. Brackenridge, J . Am. Chem. SOC.,77, 3959 (1955). (12) H. E. Ungnade, E. D. Loughran, and L. IT'. Kissinger, J . Phgs. Chem., 64, 1410 (1960). (13) Because of solvent absorption only a small portion of this band can be observed: cf. also ref. 11. (14) The somewhat higher intensity a t the maximum for the chloroethyl ether is ascribed to the effect of the chlorine which finds a parallel in the case of carbon tetrachloride.11
Volume 68, Number 11 November, 3 064
H. E. USGKADE, E. M. ROBERTS, AND
3228
minimum. This behavior is similar to that of the ethers but the interaction is weaker in this case, as would be expected.
Table I11 : Ultraviolet Absorption Bands in Aliphatic Nitro Compounds Compound
MeNOz MeNO, MeNOz MeNOz MeNOz MeNO, MeNOz MeNO, MeKOz MeNOz MeNOz MeNOz MeKOz EtNOs PrNOz BuNOz MeNOz EtNOz PrNOz BuNOz
Solvent
1MeCX HzO TL-C~F~CHZOH MeOH EtOH n-PrOH i-PrOH n-BuOH t-BuOH (ClCHzCHz)20 Et20 (MeOCHzCH2)20 Dioxane MeOH MeOH MeOH CeH14 CsH14 CBHl4 C8H14
Amin
246 247.5 246 247.5 248 248 248 248 248 248 250 252,5 265 248 248 248 246 246 246 246
6
7.0 9.2 8.3 8.5 9.7 9.5 9.5 9.8 9.7 9.6 10.4 14,O 22.6 9.9 9.1 9.4 6.5 6.8 6.8 7.1
Ampx
e
272 269 268 272.5 272.5 273 273 273.5 273.5 276 275 276 272 275 275 276 278 278 279 279
15.0 15.3 15.4 15.9 16.9 17.3 17.5 17.9 18.2 19.5 17.1 20.2 22.9 18.9 20.1 23.7 16.9 19.4 22.1 23.0
I n the case of alcohol solutions the band intensities are in the order n-C3F7CH20H < MeOH < EtOH < n-PrOH < i-PrOH < n-BuOH < t-BuOH, which corresponds to the electron-donor properties of these solvents. I n the related carbonyl absorption band, the formation of hydrogen bonds causes an increase in intensity,lE i.e., the solvent order is reversed, so that such bonding cannot account for the observed results if the solvation processes are strictly comparable. 16,17 It is proposed, therefore, that alcohols can act also as electron donors toward nitro compounds. Further evidence for such interaction comes from a consideration of absorption at the minimum. Hydrocarbons and heptafluorobutanol, which are poor electron donors, have nearly identical absorption below the minimum while the other alcohols' ROH all show an apparent red shift of 1.5-2.0 mp of the highintensity band and an increased absorption a t the
The Journal of Physical Chemistry
L.W. KISSIR'GER
Experimental Infrared Absorption Spectra. Carefully weighed amounts of alcohols and nitro compounds were made up to exact volumes, and the solutions were run in matched sodium chloride cells in a Perkin-Elmer Mod 421 infrared' spectrophotometer. For the determination of absorbance each solution was run two or three times under the same conditions. The listed intensity values represent the average. Ultraviolet Absorption Spectra. The observed differences in intensities and wave lengths were smaller than differences between instruments and it was, therefore, not possible to use literature data. The purest available specimens of nitro compounds (99.9+%) were weighed out into volumetric flasks and dissolved in Spectrograde solvents, Absorption spectra were determined in silica cells 0.2 cm. thick (to minimize solvent absorption) with a Cary Model 14 spectrophotometer a t 25'. The intensities in the present investigation were reproducible to within 0.5%. N.m.r. Spectra. Freshly prepared solutions with the same concentrations as those used for the infrared determinations were run within 1 hr. on a T'arian DP-60 spectrometer. The chemical shifts were measured by the side-band technique. I n 0.5 f W methanol solution cyclohexane was used as external reference while in 0.1 M methanol solutions tetramethylsilane was used as internal reference. All data in Fig. 2, however, are reported with reference to the methyl peak of methanol. The temperature of the probe was 30'. (15) C. N. R. Rao, G. K. Goldman, and A. Balasubramanian, Can. J . Chem., 38, 2508 (1960). (16) N. S. Bayliss and E. G. McRae, J . Phys. Chem., 5 8 , 1002, 1006 (1954). (17) Polarization effects which depend on the solvent refractive index are probably obscured since they would lead t o the order n-Bu > t-Bu, opposite t o t,he experimental findings.
