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Intermolecular and intramolecular hydrogenbonding in 8-quinolinol and 2-methyl-8-quinolinol
Item Type text; Thesis-Reproduction (electronic)
Authors Swartz, William Richard, 1945-
Publisher The University of Arizona.
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Download date 13/06/2018 13:34:33
Link to Item http://hdl.handle.net/10150/347688
«I1TERM0LECULAR AKD- II'ITRAJ'IOLECUMR HY'DROGEN BOBDIHG IE
• 8-QUIEOLIUOL M D 2-METHiL-8-QUIEOLIIOL
"byWilliam Richard Swartz
A Thesis Submitted to the Faculty of the> - "
DEPARTMENT OF CHEMISTRYIn Partial Fulfillment of the Requirements
For the Degree of'MASTER OF SCIENCE
In the Graduate CollegeTHE UNIVERSITY.OF ARIZONA '
1 9 6 9
STATEMENT BY AUTHOR
This thesis has been submitted in partial fulfillment of requirements for an advanced degree at The University of Arizona and is deposited in the University Library to be made available to borrowers under rules of the Library.
Brief quotations from this thesis are allowable without special permission, provided that accurate acknowledgment of source is made. Requests for permission for extended quotation from or reproduction of this manuscript in whole or in part may be granted by the head of the major department or the Dean of the Graduate College when in his judgment the proposed use of the material is in the interests of scholarship. In all other instances, however, permission must be obtained from the author.
SIGNED
APPROVAL BY THESIS DIRECTOR This thesis has been approved on the date shown below:
<r <- U
QUINTUS FERNANDO y DateProfessor of Chemistry
TO Mr MOTHER
ACKNOWIEDGMEajTS
I would like to thank. Dr. Quintus Fernando for his guidance throughout the course of this investigation.
TABLE OF CONTENTSPage
LIST' OF TABLES ............ . v±LIST OF ILLUSTRATIONS . ........... • ■ viiABSTRACT ......... L . ................. viiiINTRODUCTION . . . . . . . ... . . . . . . . . . . . . . . 1
Hydrogen Bond . . . . . . . . . . . 1Intramolecular Hydrogen Bonding . . ................. 2,Intermolecular Hydrogen Bonding .............. 5Nuclear Magnetic Resonance Spectrometry . . . . . . . . . 8Comparison of the Results Obtained from Infrared to -
Nuclear Magnetic Resonance Techniques . . 12Statement of Prdblem-,. . . . . . . . . . . . I . . . . . . 13
EXPERIMENTAL . . . «'.......... ................ .......... 'lkX -
Materials ............ . . . . . . . . . l4Infrared Spectrophotometry . . . . . . . . . . l4Nuclear Magnetic Resonance Spectrometry . . . . . . . . . l4
RESULTS ................. 16Nuclear Magnetic Resonance Data . . . . . . . . . . . . . l6Infrared Spectrophotometric Results . . . . . ............ 28
DISCUSSION . . . ................... 37, Infrared Data . . . . . . . . . . . ................... 37.Nuclear Magnetic Resonance Interpretation . . . . . . . . 38Proposed Equilibrium...... ................... 4l
REFERENCES................. . . 43
v
LIST OF TABLESTable Page1. Limiting Hydrogen Bond Distances "between Oxygen and
Nitrogen Atoms . . . . . . . . . . . . . . . ........ . 12. Energy of Hydrogen Bond Formation with Oxygen and
Nitrogen Atoms . . . . . . . . . . . . . ........ . . . . 23* Classes.of Non Hydrogen Bonding Solvents . . . . . . . . 44. Chemical Shift of the Hydroxyl Resonance in 8-quinolinol
in Chloroform ................. 245 . Chemical Shift of the Hydroxyl Resonance in 2-methyl-
8-quinolinol in Chloroform ............. 25
6. Chemical Shift of the Hydroxyl Resonance in 8-quinolinolin Carbon Tetrachloride . . . . . . . . . . . . .......... 26
7- Chemical Shift of the Hydroxyl Resonance in 2-methy1-8~quinolinol in Carbon Tetrachloride . . . . . . . . . . 27
8 . Change in the Infrared Absorbance and Half Width ofthe Hydroxyl Band in 8-quinolinol in Chloroform . . . . 29
9 . Change in the Infrared Absorbance and Half Width of the Hydroxyl Band in 2-methy1-8-quinolinol inChloroform ................. 31
10. Change in the Infrared'Absorbance and Half Width" in the Hydroxyl Band in 8-quinolinol in CarbonTetrachloride ............. . . . . . . 33
11. Change in the Infrared Absorbance and Half Width in the Hydroxyl Band in 2-methyl-8-quinolinol.inCarbon Tetrachloride . . ............................ 35
vi
LIST OF ILLUSTRATIONSFigure1. 60 MHz BMR Spectrum, of 8-quinolinol (2 .3 M) in CCl^ . . , .2. . 60 MHz HMR Spectrum of 8-quinolinol (2,5 M) in CDC1„ . . ,d .
3. 100 MHz MMR Spectrum of 8-quinolinol (2,J M) in CHC1 . . .d4. 100 MHz HMR Spectrum, of 8-quinolinol (l,2.M) in CHC1„ . » ,5 . 60 Ifflz EI'IR Spectrum of 2-methyl-8-quinolinol (2,1 M)
' in CCl^ . .6 . 100 MHz HMR Spectrum of 2-methyl-8-quinolinol (2,5 M)
in CDCl. . . . . . . . . . . . . . . . . . . . . . . . . .37 . 100 MHz HMR Spectrum of 2-»methyl-8-quinolinol (2 .7 M)
in CHC1 . . . . ........ . . . . . . . . . . . . . . . .o8. Chemical Shift of the Hydroxyl Resonance of 8-quinolinol
m CHCl^9 . Chemical Shift of the Hydroxyl Resonance of 2-methyl-
8-quinolinol in CHCl^ . . . . . . . . . . . . . ............
ABSTRACT
Intramolecular and intermdlecular hydrogen "bonding equilibria in 8-quinolinbl and 2~methyl-8-quinolihol were studied by infrared spectrophotometry and nuclear magnetic resonance spectrometry. While nuclear magnetic resonance spectrometry showed that an equilibrium, exist between associated and non-associated 8-quinolinol and 2-methyl- 8-quinolinol, infrared spectrophotometry could not detect these species that were present at equilibrium.
ylii
INTRODUCTION
In this work hydrogen bond formation between oxygen and nitrogen atoms in heterocyclic organic molecules will be investigated. The various experimental methods which have been employed to study hydrogen bond formation will be surveyed with an emphasis on infrared and nuclear magnetic resonance techniques.
Hydrogen BondIf the bond distance between two atoms A and B in the bond A-H • • -B
is shorter than the van der Waals distance between A and B, then hydrogen bonding can be said to exist. On the other hand, if the distance between atoms A and B is large enough to be explained by van der Waals forces then hydrogen bonding does not exist. In general, if the distance between A and B, in A-H • • *B is greater than 3.4 Angstroms there is no hydrogen bonding between atoms A and B. The limit of the bond distance in Angstroms for which hydrogen bonding can be considered to occur when A and B are nitrogen and oxygen atoms is listed in Table 1 (Pimentel i9 6 0).
Table 1Limiting Hydrogen Bond Distances Between Oxygen and Nitrogen Atoms
Limiting Bond Distance
Type of Bond
3.2 A 0-H - • -03 .5 A N-H*•*N3.3 A 0-H-•-N3.4 A N—H - - * 0
1
2
In Table 2 are listed the enthalpy values necessary for the formation of several specific hydrogen bonds (Pimentel i9 6 0).
Table 2
Energy of Hydrogen Bond Formation with Oxygen and Nitrogen Atoms
Type of Bond H
0-H-•-0 3 -6 kcal/moleN-H* *-N 3 -5 kcal/moleN —H • • • 0 4.2 kcal/mole
Operational criteria for hydrogen bond formation are quite numerous. Some of the more important criteria for hydrogen bond formation are the changes in the frequency of absorption, increase in band width and increase in intensity in the infrared spectra; a shift toward lower field in the nuclear magnetic resonance spectra, and a molecular weight which is larger than a, formula weight as shown by cryoscopy, vapor pressure and vapor density data are other indications of hydrogen bonding. Further criteria are solubility studies, dielectric constant, partition or distribution, molar volume, parachor, refractive index, electrical thermal conductivity, Raman spectra, neutron and x-ray diffraction studies. The experimental techniques that will be employed in this study will be infrared and nuclear magnetic resonance spectroscopy.
Intramolecular Hydrogen Bonding Hydrogen bonding can be classified either as intramolecular hydrogen
bonding or intermolecular hydrogen bonding. An intramolecular hydrogen bond is a hydrogen bond formed in one molecule while intermolecular hydrogen
bonding is hydrogen bonding that occurs between at least two molecules.The following conditions are necessary for intramolecular hydrogen bonding: (l) the bonds can form only under specific spatial requirements,(2) the formation of the bonds does not result in chains of molecules.Since intramolecular hydrogen bonds must be studied in solutions of low concentrations, nuclear magnetic resonance is of little use in studying this type of bond. On the other hand, infrared spectroscopy is sensitive to hydrogen bonding over a wide concentration range.
Successive dilution makes it possible to study monomers by eliminating intermolecular hydrogen bonding. A free non-assbciated hydroxyl group in an alcohol or a. phenol produces a band at 3600-3650 cm. "** (Hadzi 1959)* The free hydroxyl group' is affected by its steric environment. For instance, in cyclohexanols the group frequencies of the axial hydroxyl groups are higher than those of the equatorial hydroxyl groups. Since free hydroxyl groups have different properties than bonded hydroxyl groups equation.(l) has been proposed to represent this difference (Tichy 1965).
Av(OH) = v(OH)f -v(QH)Av(OH) = the change in frequency • ' ( ’O 'v(OH)̂ , = frequency of the free hydroxyl group v(OH)^ = frequency of the bonded hydroxyl group
Bonded hydroxyl groups have increased intensities, greater half widths and higher molar absorptiviti.es in the infrared. It is, however> not possible to measure the concentrations of the associated and .nonassociated species from the band intensity, band width and molar absorptivity. On the other hand, a, rough estimate can be made of the
concentration of associated and non-associatcd species by using these band characteristics.
Conditions of MeasurementsSolvents such as carbon tetrachloride and carbon disulfide arc
usually suitable for hydrogen bonding studies since they do not absorb in the 3 P region and since they do not, in general, form hydrogen bonds with the solutes. Pimentel (195^) has listed some solvents with respect to their hydrogen bonding ability which have been compiled in Table 3*
Table 3Classes of Mon Hydrogen Bonding Solvents
Solvents
1. CCl̂ ; CSg2. CHjCCl ; 0H I Hydrogen Bonding
J increases3. CHBr^; CHC1 ; CHgClg
In all cases the presence of hydrogen bonding was deduced by observing the bonded MH(3220 cm or the free MH( 3̂ -50 cm "") in N-ethyl acetemide. In later studies Huggins and Pimentel (1955) observed hydrogen bonding between the solvent and deuterated chloroform by observing the increase in half width and intensity of the C-D band. They listed the following solvents in order of their decreasing hydrogen bonding properties:CCl^ < CHCl^ < 0H < 0MOg < acetone < tricthylamine. Obviously, when studying hydrogen bonding in a specific molecule an inert solvent must be used or the data will be meaningless.
5
Concentration of .005 M must be used to avoid intermolecular association. Even at this concentration, some diols exist as dimers.The only method of establishing whether a bond is intramolecular is to prove that its molar absorptivity is independent of concentration. When intermolecular bonds are formed in addition to intramolecular bonds, extrapolation of the absorptivity to zero concentration does not pass through the origin (Tichy 1965)•
The possibility of mistaking other types of bond formation for hydroxyl bonds can be eliminated by using deuterated compounds. A decrease in band intensity or disappearance of the band in question and the appearance of a band of the same characteristics in the 2400-2700 cm ^ range is proof that the original bond involved the hydroxyl group.
As a point of interest it is noted that the atoms involved in the strongest hydrogen bonds are collinear. Any deviation from collinearity will weaken the bond. Also, to this date, no reliable quantitative correlation between enthalpy and change of hydroxyl absorption is known for intramolecular hydrogen bridges, although usually a larger change of frequency corresponds to a stronger hydrogen bond.
Intermolecular Hydrogen Bonding As stated previously, intermolecular hydrogen bonding consists of
hydrogen bonds formed between at least two molecular species. Infrared • spectrophotometry readily distinguishes intermolecular hydrogen bonds from intramolecular bonds. A change in the frequency of absorption (AV ) versus the half width (AV 1/2) is large for intermolecular and small for intramolecular hydrogen bonded systems. At low concentrations, spectra
resulting from intermolecular hydrogen "bonding change radically while spectra resulting from intramolecular hydrogen bonding remain unaltered. Furthermore, varying concentration or temperature have the same effect on intermolecular bonding (Coggeshall and Saier 1951)• Since hydrogen bonding can occur with the solvent, care must be taken to distinguish solvent effects from the effects produced by the compound in question.This is usually made possible by using different solvents with various hydrogen bonding properties. , Often it is desirable to study the equilibria between the intramolecular hydrogen bonded species, the monomer, and the intermolecular bonded species, the polymer. The equilibria between these
• c
species have the following properties: (l) the hydrogen bonding systemsinvolve monomers and several polymers in rapid equilibrium, (2 ) each polymer has a characteristic frequency of absorption (V ) and the higher the polymer the lower is the V , (3 ) each polymer has a characteristic molar absorptivity and the higher the polymer the higher the molar absorptivity (4) since the molar absorptivity (a) may increase by a,s much as a factor of ten on bond formation, a small shift in equilibrium produces a large spectral change.
Hydrogen bond formation in alcohols has received a good deal ofinterest. Errera and Mollet (1936, 1940) established the spectralcharacteristics of alcohols over 30 years ago. Sharp bands in the region
_1of 3o30 cm which decreased;in intensity with increasing concentration_ ~1were attributed to the monomer. Broad bands from 3300-3500 cm which
decrease in intensity with decreasing concentration were assigned to the hydrogen' bonded species. Hydrogen bonding in benzyl alcohol was also studied (Fox and'Martin 1936). • -The results were interpreted - in terms- of
7
an equilibrium between hydrogen bonded- species- Kempter and Mecke (l94l) examined phenol and determined an equilibrium constant for the hydrogen •bonded species. They assumed that complexes of all orders were possible and that the equilibrium constants for the complexes were identical. Unfortunately, this assumption cannot be validated by .any experimental evidence. Coggeshall and Saier (1951) made a further investigation of this system. They determined a constant for the monomer and dimer equilibrium and a. constant for all the other equilibria. Once again an assumption was made, with little supporting evidence, that a dimer was formed, when the trimer or even the tetramer may be more stable if one considers a collinear hydrogen bond to confer stability on the molecule . (Coburn and C-runwald 1958). Liddel and Becker (1957) made an extensive study of the lower alcohols. By using data obtained by other experimental methods they were able to decide whether a dimer, trimer, or higher polymer were formed. With this knowledge they were able to determine equilibrium constants for some of the hydrogen bonded species.
The obvious pitfall in most of the determinations of equilibrium constants between hydrogen bonded species is that the existence of specific species has to be assumed. Another problem with this method of determination stems from a practical point. The species in question must be moderately soluble in a non-hydrogen bonding solvent. This eliminates the possibility of studying many types of systems by this method. Also, each species to be studied must have a distinctive absorption in the infrared which is proportional to its concentration. Unfortunately, most spectra overlap with or partially distort other absorption bands and the concentration of a particular species can only- be approximately obtained.
For instance, in most studies the determination of the equilibrium constant is almost always based on the intensity measurements of the absorption band of the monomeric species and the assumption that the monomer concentration is proportional to the absorption at this frequency. This assumption obviously becomes invalid if any other species absorb at this frequency. ,
Nuclear Magnetic Resonance SpectrometryArnold and Packard (1951) found that the chemical shift of the
hydroxyl proton in ethanol was concentration and temperature dependent. Liddel and Ramsey (1951)' stated that the temperature dependence of the resonance signals could be understood if there were alternative molecular states whose energy of separation were of the order of KT. Since ethanol forms hydrogen bonds involving the OH group, the hydrogen in the hydroxyl group should experience different magnetic shielding in associated and unassociated states. If the correlation time, the time related to the lifetimes of the two states, is small (one millisecond) resonance will be observed at the frequency corresponding to the average shielding of the two states. Thus, changes of temperature and concentration alter the associated and unassociated states, and resonance frequencies will be temperature and concentration dependent.
Cohen and Reid (1956) used solvent dilution methods to obtain hydrogen bond shifts for methanol and ethanol in various solvents. On extreme dilution the hydroxyl peak shifted to a higher field. These results can be explained by considering the bond A-H-•-B where B is the. donor atom. The associated bond is 'shifted to a lower field due to the
following effects: (l) The donor atom tends to draw protons away fromthe bonding electrons and reduces the electron density in the immediate vicinity of the hydrogen atom. (2 ) The diamagnetic circulations within the hydrogen atom is inhibited by the presence of a strong electric field.
Shifting of associated bonds to a higher field was also confirmed by Korinek and Schneider (1957)• They measured the shift of the proton resonance signal of chloroform in solvents ranging from hexane, an inert solvent, to triethylaiaine, a donor solvent, and found that an increase in association resulted in a shift to lower field. In practice the equilibria between associated and non~associated species are studied by diluting the species in question with an inert solvent. Concentrations as low as .1 M and temperatures varying from -80 to 200° C. can be used to study hydrogen bonding .equilibria. From these studies the chemical shift, which is a measure of the electron density around the proton, and the exchange times, which are a measure of the times it takes for a proton to go from an associated to a non-associated state, can be obtained. •
Several studies of associated and non-associated species have, been made for phenols and lower alcohols. Bartdorf (1955) obtained different shifts for phenols depending on which solvents he used. Cyclohexane, an inert solvent, produced a slight shift to higher field on dilution.Dilution of phenol with acetone to 60 mole percent gave rise to a low field shift. This implies that a stronger bond is formed between acetone and phenol than between the associated phenol molecules. The OH signal for ortho methyl substituted phenol in the pure compound occurs at a higher field than the signal of pure phenol itself. Also, the shift upon dilution.of. the ortho substituted-phenol in acetone was found to.be .
larger than that of phenol in acetone. This indicates that the pure ortho substituted phenol has a weaker hydrogen bond than that of pure phenol. Undoubtedly this phenomenon is due to the steric hindrance caused by the methyl group in the pure compound. In agreement with this* pure 2,6-diisopropyl phenol, which has a greater steric effect, has an OH peak at a higher field than any of the other phenols. It was also noted that the hydroxyl peak from pure ortho nitro and carbonyl substituted phenols occurh at a lower field than the pure methyl and isopropyl substituted phenol. Since ortho nitro and methyl substituted carbonyl compounds form strong intramolecular hydrogen bonds, there is comparatively little association occurring in these compounds at higher concentrations in comparison with the ortho methyl substituted compounds. Further studies by Huggins, Pimentel and Shoolery (1956) showed that at low concentrations there was no change in the chemical, shift for ortho chlorophenol as compared to raeta and para ohlofophenol. ■ It can thus be concluded that at low concentrations ortho chlorophenol is stabilized by intramolecular hydrogen bonding.
In one of the better studies Becker, Liddel and Shoolery (1958) used infrared measurements to show the fraction of the hydroxyl proton of ethanol that was.not hydrogen bonded. This method permits•subtraction of the chemical shift of the unbonded ,species from tlie overall chemical shift. The remainder of the chemical shift is an average which is determined by all the other'hydrogen bonded species. If another technique is available to show at which concentrations monomers and dimers only are present it can justifiably be assumed, that certain portions of the nuclear magnetic resonance spectrum are a result strictly of'equilibrium between
11
monomers and dimers. Unfortunately most studies simply assume that this equilibrium exists at low concentrations without sufficient evidence to support this assumption. Even Becker assumes monomer-dimer equilibria when monomer-trimer equilibria might be more favorable as shown by Coburn (1958).
In many of the equilibrium studies of hydrogen bonded species the following assumptions are made: (l) at extremely low concentrations onlythe monomer exists, (2) at high concentrations the chemical shift is primarily determined by monomer-polymer equilibria, (3 ) at relatively low concentrations the chemical shift is influenced strictly by the monomer-dimer equilibzdum. These assumptions can be subject to a great deal of error. In a detailed study of ethanol Decker, Liddel and Shoolery (1958) examined the spectra at very low concentrations (.005 M). The chemical shift remained, constant below .01 M. This is an indication that only the monomer exists below .01 M and that above this concentration the chemical shift is influenced by other hydrogen bonded species beside the monomer. In most investigations, concentrations below .1 M cannot be studied in the nuclear magnetic resonance frequency range. In this case the chemical shift due to only the monomer is usually obtained, by extrapolating low concentration values to infinite dilution and assuming ' that this extrapolation gives the chemical shift which is related -to:only monomer concentrations (Huggins, Pimentel and Shoolery 1956). This assumption may be invalid. If there are no other data which can furnish information about the equilibria of hydrogen bonded species at low concentrations the assumption should not be made.
12
Comparison of the Results Obtained from Infraredto
Unclear Magnetic Resonance TechniquesIn the infrared region vibrational bonds for the associated
and non-associated hydrogen bonded species can be studied simultaneously.This is not possible in nuclear magnetic resonance method due to largedifferences in the frequencies of the two techniques. To observe bothassociated and non-associated hydrogen bonded species in the same medium,nuclear .magnetic resonance measurements would require lifetimes of eachstate to be longer than the reciprocal of the hydrogen bond shift, expressed
-2in frequency units. The reciprocal of the shift may have values of 10 to 10 seconds. Since only one sharp signal is observed, the lifetime of the hydrogen bonded states must be shorter than the time required for detection. Thus in this amount of time, 10 to 10 seconds, the hydrogen bond will be broken and reformed many times.
Nuclear magnetic resonance is sensitive to rather weak interactions which do not appear in the infrared. For stronger interactions both methods are complementary. The infrared method is superior for studying interactions at low concentrations; although this situation can be simulated in the nuclear magnetic resonance method by operating at higher temperatures when higher concentrations are used. Although most exchange times that occur in hydrogen bonded species occur too fast to be detected, NMR can be used to measure slower exchange rates. By lowering the temperature or increasing the viscosity of the solvent the exchange rates can be slowed to a point where they can be studied. For example,-Mays and Brady (1956) measured the behavior of water adsorbed on TiQg and found that water becomes, immobilized when the coverage .
13■ . '
reaches a monolayer. Hichmott and Selwood (1956) measured the proton relaxation of alcohols and water adsorbed on catalytic solids.
Statement of .Problem 8-quinolinol and 2-metbyl-8-'quinolinol exhibit different metal
complexing properties. Recently the nuclear magnetic resonance spectrum of 8-quinolinol and 2-methyl-8-quinolinol have been reported by Hirsh and Abarbaro (1968) and Baker and Sawyer (3.968). In these studies a signal attributable to the hydroxyl proton was not observed. In the following the hydroxyl peaks in the nuclear magnetic resonance spectra of 8-quinolinol and 2-methyl-8-quinolinol will be. identified and the effect of concentration on the spectra will be studied. An attempt, will be made to interpret these results on the basis of hydrogen bonding and confirmatory evidence from infrared studies will be sought.
EKPERIMEMTAL
MaterialsAll materials used were of. analytical reagent grade quality.
8-quinolinol and 2-methy 1-8-quinoline1 were sublimed, recrystallized from 95% ethanol and dried in a vacuum desiccator. The solvents used in the infrared and nuclear magnetic resonance study were chloroform
oand carbon tetrachloride. These solvents were stored over 4a molecular sieves to prevent contamination by water.
Infrared Spectrophotometry \
Infrared, measurements were made with a 337# grating spectro-7 photometer (Perkin-Elmer) from 4000 to 3000 cm~\ Samples were contained in a sealed sodium chloride cell. Spacers were used to obtain optimum cell widths for all the concentrations examined. The spacers were .01l4, .0254, .0368 and .099 cm. in thickness. Spectra from 4000 to 650 cm were recorded on the Perkin-Elmer Infracord spectrophotometer.
v
Huclear Magnetic Resonance SpectrometryNuclear magnetic resonance spectra were obtained with a. Varian
Associates Model A -60 as well as their HA-100 spectrophotometer. A sample tube was fitted with a coaxial capillary tube in which the standard was placed. The same sample tube was used in each instance to increase the overall accuracy of the measurements. Tetramethylsilane (TMS) was used as an external standard with the HA-100 instrument. The .sweep width of .50 cycles, per second (ops) .gave an accuracy of - .0 1 cps.
Ik
However, small changes in the temperature of the probe made the re™ suits reproducible to only ~ 1 cps. * Chloroform was used as a solvent for the data obtained with the HA-100.
Spectra were recorded at $00 cps on the A~6o instrument and nitromethane was used as an external standard.' . This particular standard was used so that the peak from the standard and the spectrum of the compound in question could be recorded concurrently without making any - adjustments in the instrumental parameters. Even with this precaution the field was unstable and the information obtained can only be qualitative in nature. Carbon tetrachloride was used as the. solvent for the spectra that were obtained with the A~6o instruments
RESULTS
Nuclear Magnetic Resonance DataSpectra
Previously, Hirsh and A'barbaro (1968) and Baker and Sawyer(1968) have reported the spectra of S-quinolinol and 2-methyl-8~quinolinol without any mention of the hydroxyl peak.. For thisreason the spectra of 8-quinolinol and 2-methyl~8-quinolinol will he
*6 .
listed along with the identification of the hydroxyl peak.
Figure #1: 60 MHz NMR spectrum of 8-quinolinol (2 .3 M)in CCl^,.hydroxyl peak at 9-50 cP
Figure //2: 60 liHz NMR spectrum, of 8-quinolinol (2 . 5 M)in CDCl^, hydroxyl pea.k at 9 ”25 cf1
Figure #3: 100 MHz NMR spectrum of 8-quinolinol (2 . 7 M)■ in CHCly hydroxyl peak at 9-14 cp
Figure $k: 100 MHz NMR spectrum of 8-quinolinol (1.2 M)in CHCl^, hydroxyl peak disappears at lowconcentration
Figure $5? 60 MHz NMR spectrum of 2~methyl-8-quinolinol(2 . 1 M) in CCl^, hydroxyl peak at 8.93 cT
Figure ^6 : 100 MHz NMR spectrum of 2-methyl-8-quinolinol(2 . 5 M) in CDClg, hydroxyl peak at 8 . 6 2 cf
Figure ^7: 100 MHz NMR spectrum of 2-methyl-8~quinolinol(2 . 7 M) in CHC1q, hydroxyl peak at 8.62 cf j
-I.10 9.2b 9 8 P.'M I $)
Figure 1: 60 MHz MI-IR spectrum of 8-quinolinol (2.3 M) in CC1,
8 PPM (5)
Figure 2: 60 MHz KMR spectrum of 8-quinolinol (2.5 M) in CDCl^
18
, pTr
10 9.-1
> ”>
8 7 jJ' " 6 " ‘ S PPM iS) 4' ""j
Figure 3= 100 MHz IJI'IR spectrum of 8-quinolinol (2 .7 M) in CHCl^
Figure b: 100 MHz HMR spectrum of 8-quinolinol (1.2 M) in CHC10
19
PPM IS) 7
Figure 5'« 60 MHz IJI'-iR spectrum of 2-methyl-8-quinolinol (2.1 M) in CCl^
,<VWtV\
8 6?
Figure 6: 100 MHz I'IMR spectrum of 2-methyl-8-quinolinol (2.5 M) in CDC10
20
5 PPM 18)
Figure 7: 100 MHz MMR spectrum of 2-methyl-8-quinolinol (2 .7 M) in CHCl^
A plot of the chemical shift of the.hydroxyl resonance in 8-quinolinol and 2-methyl-8-quinolinol versus the molar concentration in chloroform are shown in Figures 8 and 9 respectively. Tables of the chemical shifts of the hydroxyl resonance versus concentration for 8-quinolinol and 2-methyl-8-quinolinol'in chloroform and carbon tetrachloride are shown in Tables 5̂ 6 and 7 - The‘data in Tables 6 and 7 were recorded on the Varian A-6o spectrophotometer and are intended for qualitative use only.
22
9.525
9.396
9.267
89.138
9.009
- V
8.880 L- ° 1.2
.©
q A d
O'Zo
o
1.65 2.10 2 .55 3.00
CONCENTRATION (M)
Figure 8: Chemical Shift of the Hydroxyl Resonance in 8-quinolinol CHC10
23
CONCENTRATION
Figure $): Chemical Shift of the Hydroxyl Resonance in 2-methyl-8-quinolinolCHC13
2k
Table 4Chemical Shift of the Hydroxyl Resonance
in8-quinolinol in Chloroform
Instrument: HA-100Sweep width: 50 cps
o § o ES
Position of Hydroxyl Resonance ((f') V 1 /2 (cps)
2 .9 0 9 9.524 6.32 .7 8 0 9.490 7.42 .6 9 0 9.473 7.22 .5 9 5 9.429 7.82.481 9.427 92.373 9.385 9.22 .2 6 5 9.3^7 10
2.157 9.306 1 2 .5
2 .0 5 0 9.277 13
I .618 9 .1 2 16
1.510 9 .0 0 17i.4oo 8 .94 191 .2 9 0 8 .8 8 30
1 .187
25
Table 5Chemical Shift of the Hydroixyl Resonance
in2~methyl-8-quinolinol in Chloroform
Instrument: HA-100Sweep width: 50 cps
Cone, (M)Position of
Hydroxyl Resonance ((p) V 1/2 (cps)
3.170 8.735 2.93.003 8 .7 1 8 2.92 .7 8 9 8 .6 8 3 3.02 .6 7 8 8 .6 8 8 3.02 .9 6 6 8 .6 6 9 3.02.455 8 .6 5 1 3.82.343 8 .6 3 3 4 .0
2 .2 3 1 8 .6 1 6 4 .1
2 .1 2 0 8 .5 8 5 4.42 .0 0 8 8 .5 6 6 4.4
1.897 8.540 5-0
1.785 8 .5 1 2 5-0
1.673 8 .4 7 7 5 .6
1 .562 8 .4 5 6 5.61.450 8.403 6 .6
1.339 ---- ; - ; — • - * •
t
Table 6
Chemical Shift of the Hydrovtyi Resonancein
8-quinolinol in Carbon TetrachlorideInstrument: A~6oSweep width: ^00 cps
Cone. (M)Position of
Hydroxyl Resonance (if)
2.3% 9.%982 .2 5 9.4522 .1 5 9.4402 .0 6 9.415
1.97 9.415I .87 9.4101 .6 8 9.390
Table 7Chemical Shift of the Hydroxyl Resonance
in2-methyl-8-quinolinol in Carbon Tetrachloride
Instrument: A-60 Sweep width: 500 cps
Cone. (M)Position of
Hydroxyl Resonance (cT)
2 .2 6 8.9152 .1 7 8 .9 0 7
2 .0 8 8 .8 9 9
1.99 8 .8 9 0
1.90 8 .8 7 3
1 .8 0 8 .8 6 5
1 .7 2 8 .8 3 2
1 .6 3 8.8%91.53 8 .8 0 6
1.45 8 .8 2 9
1.13 8.7581 .0 2 8 .7 4 1
28
Infrared Spectrophotometric Results
A band at 3^00 cm ^ in the infrared region was identified as the band due to the hydroxyl groups of the 8-quiholinol and of the 2-methyl-8-quinolinol through dilution experiments. This band did not shift frequencies at various concentrations; although the band did increase in half width ("V" 1/2 y at higher concentrations <■In Tables 8, 10 and 11 the absorbances (A), concentrations andhalf widths are reported for 8-quinolinol and 2-inethyl-8-quinolinol in chloroform and. carbon tetrachloride.
29
Table 8Change in the Infrared Absorbance and Half Width
of theHydroxyl Band in 8-quinolinol in Chloroform
Cell •width: 0.1143 cm
Cone. (M) A V 1/2 (cm"1)
2 .1+81 .3468 2592.373 .3426 2442 .2 6 5 .3233 2332 .1 5 1 .3279 211
2 .0 5 0 .3233 211
1.942 .3054 2051.834 .3054 199I .726 - .2840 1941 .618 .2798 171
1.51 .2798 178
l.4o .2557 1711 .187 .2291 1711 .0 7 9 .1904 142
.9708 .2111 138
.8629 .2041 144
.755 .1643 116
.647 .1487 116
.539 .1308 116
.431 .1079 ill
.323 .0888 ill
.216 .0580 105
.130 . 0348 100
.108 .0292 100
Table 8 — Cont’dCell width: .0368 cm
Cone. (M) Cone, x 3*2 (M) V 1/2 (cm”1)
.027 .09 87
.216 •71 111
.434 1.39 144
.647 2 .0 9 178
Cell width: ,02^4 cm
Cone. (M) Cone, x 2.2 (M) V 1/2 (cm”1)
.027 . 06 100
.2157 .48 10$
.434 .9 6 133
.647 1.44 155
31
Table 9Change in the Infrared Absorbance and Half Width
of theHydroxyl Band in 2-methyl-8-quinolinol in Chloroform
Cell width: .011^3 cm
Cone. (M) A V 1/2 (enf1)
3.033 .6021 167
2 .7 8 9 .6021 156
2 .6 7 8 .5850 1562.1&55 .5768 1442 .2 3 1 .5452 1392 .0 0 8 .5086 1331.897 .5086 1331.785 .5058 1331 .562 .4437 122
1.339 .3979 122
1.129 .3420 111
.924 .2882 105
.755 .2441 100
.540 .1805 100
.328 .1135 100
.108 .0434 89
32
Table 9 — Cont’dCell width: .0368 cm
Cone. (M) Cone, x 3.2 (M) V 1/2 (cm”1)
.540 1.73 133
.324 l.o4 115
.107 .348 100
Cell width: .0254
Cone. (M) Cone, x 2 .2 (m) V 1/2 (cm”1)
.924 2 .0 3 139
.755 1 .6 6 133
.539 1 .18 122
33
Table 10Change in the Infrared Absorbance and Half Width
in theHydroxyl Band in 8-quinolinol in Carbon Tetrachloride
Cell width: .011^3 cm
Cone. (M) A V 1/2 (cmT^)
2 .15U .7328 178
I.#? .6882 144
1.873 . 6716 144
1.777 .6925 1441 .6 8 6 . 6478 1391.405 .5768 122
I .171 .5157 1171.054 .4750 105
.820 .3979 100
.586 .2882 78
.351 .2041 67
.i4o .0888 61
.070 .0482 55
31!
Table 10 — ContfdCell width: .0368 cm
Cone. (M) Cone, x 3*2 (M) V 1/2 (cm”1)
.585 1 .8 9 133
.351 1.13 100
.l4l .45 66
Cell width: .02^h cm
Cone. (M) Cone, x 2.2 (m ) V 1/2 (cm""1)
0OJCO 1 .8 0 133.586 1 .2 9 100
.351 •77 88
35
Table 11Change in the Infrared Absorbance and Half Width
in theHydroxyl Band in 2-methy1-8-quinolinol in Carbon Tetrachloride
Cell ■width: .01143 cm
Cone. (M) A V 1/2 (cm*-1)
2 .0 7 9 .7447 122
1 .898 .7570 111
1 .717 .7100 1111 .627 .6778 1051 . 5J40 .6383 100
1 .447 .6289 92
1 .3 0 1 .6021 911 .2 3 0 .5850 891 .132 .5452 881 .0 1 9 .4949 80
.792 .4l68 78, 566 .3188 67
.340 .2041 67
.113 .0757 67
.045 .0315 55
36
Table 11 Cont'dCell width: .0368 cm
Cone. (M) Cone, x 3*2 (M) V 1/2 (cm"1)
.5660 1 .8 1 1 100
.3400 1 .088 89
.1350 .432 67
Cell width: .0254 cm
Cone. (M) Cone, x 2 .2 (m) V 1/2 (cm"1)
.7924 1.74 111
.5660 1 .2 5 89
.3396 •75 78
DISCUSSION
The data obtained from infrared and nuclear magnetic resonance studies will be compared with other data that have been published on related systems. Finally, suggestions will be proposed for further studies on hydrogen bonded equilibria.
Infrared DataA hydrogen bonding system involves monomers and several polymers
in rapid equilibrium. Each species should have a characteristic frequency of absorption in the infrared. At low concentration, spectra resulting from intermolecular hydrogen bonding change radically with concentration while intramolecular hydrogen bonded species remain unaltered. In a study of alcohols, Errera and Mollet (1936) have shown
-1that the band at 3630 cm which decreases in intensity with increasing concentration is proportional to the monomer concentration; while the band at 3300-3500 cm which decreases in intensity with decreasing concentration'is proportional to the concentration of the polymer. In this work, the infrared spectra of 8-quinolinol and 2-methyl-8-quinolinol showed a broad band, at 3^00-3500 cm The band frequency does not shift, with a change in concentration, although the half width (V 1/2) does increase with an increase in concentration (Tables 8, 9; 10 and 11). In the low concentration range (.1-.005 M) the characteristics of the band remain unchanged.
37
■ From this infrared information the following deductions can be made: (l) the band at 3400-3500 cm ^ corresponds to one speciessince it does not shift in frequency at different concentrations,(2 ) since the band is broad and does not change its characteristics at low concentrations it is representative of an intramolecular hydrogen bonded species (monomer),, (3 ) since the band width increases at higher concentrations.
Two solvents were used to distinguish the solvent effects from the effects produced by the compound in question. The solvents used in, this work were chloroform and carbon tetrachloride, Since chloroform is a stronger hydrogen bonding solvent than carbon tetrachloride, increased hydrogen bonding between the solvent and the solute should be expected for chloroform. This can be seen (Tables 8-11) from the larger half widths of the 8-quinolinol and 2-methyl-8-quinolinol bands in chloroform compared to - carbon tetrachloride.
Nuclear Magnetic Resonance Interpretation Hydrogen bonds which are not detected in the infrared region are
usually readily observable with the nuclear magnetic resonance technique. The hydrogen bonded hydrogen in the hydroxyl group should experience different magnetic shielding in associated and non-associated states. In most instances the observed-resonance frequency is a weighted average of the resonance frequencies of the very slow exchanging states. A change in temperature or concentration should alter the relative concentrations. of the associated.and non-associated states. • Therefore, the resonance
frequencies of the associated and non-associated states will be concentration and temperature dependent. Cohen and Eeid (19 5 6) used solvent dilution methods to obtain hydrogen bonded shifts for methanol, and ethanol in various solvents. In their work, at extreme dilution, the peak due to the hydroxyl resonance shifted to a higher field.Korinek and Schneider (1957) confirmed that associated bonds shift to higher fields by measuring the shifts of the proton resonance signal of chloroform in solvents ranging from-hexane, an inert solvent, to tri- ethylamine, a basic solvent. The shifting of hydrogen bonded species to a higher field can be explained by the following: (l) donor atoms tendto draw protons away from the bonding electrons and reduce the electron density in the immediate vicinity of the hydrogen atom, (2 ) diamagnetic circulations within the hydrogen atom are inhibited by the presence of a strong electric field. In a study of phenols, Bartdorf (1965) found that the resonance signal from the hydroxyl proton occurs at a higher field for pure 2-methyl phenol than for pure phenol. This effect was attributed to steric hindrance. In agreement with this, 2,6-diisopropyl phenol, which would exhibit steric hindrance in intermolechlar hydrogen bonded species, has a hydroxyl resonance frequency at a higher field than. . either the phenol or the ortho substituted phenol.
From the above information and the nuclear magnetic resonance studies on 8-quinolinol and 2-methyl-8-quinolinol (Figures 1, 2 and Tables 6, 7 ) the following can be deduced: (l) the resonance signal that,is affected by a change in concentration (disappears at low concentrations) is representative of the hydroxyl proton, (2) since the frequency of this • resonance signal is•very concentration dependent, an equilibrium between
4o
hydrogen bonded species is being established, (3 ) the hydrogen bond responsible for equilibrium between the associated and non-associated species is comparatively weak since it was not located in the infrared spectra, (4) since the peak shifts to a lower field at higher concentration an equilibrium is established between an associated hydrogen • bonded form and a non-associated hydrogen bonded form (monomer), (5 ) the chemical shifts in 2-methyl-8-quinolinol occur at a higher field than the chemical shifts of 8-quinolinol.which are due to a, steric effect, an inductive effect or both.
The exchange time is a measure of the time it takes for a protonto go from an associated to a non-associated state. This time is inversely proportional to the band width (V 1/2) in the nmr; that is, thelarger the band width the shorter the exchange time. In Tables 4 and5 the increase in the band width with decreasing concentration indicates that as the concentration of the solute increases the hydrogen is exchanging less rapidly.
To distinguish between the hydrogen bonding in the solute from the hydrogen, bonding with the solvent, chloroform and carbon tetrachloride were used. It.can be seen from Tables 6, T, 8 and 9 that the. chemical shift in chloroform is twice that of the chemical shift in carbon tetrad- chloride. Thus, the equilibrium, in the associated and non-associated states in chloroform is affected to a slight extent by. the hydrogen bonding properties of chloroform.
Proposed equilibrim.In the infrared spectra only one absorption band that is
attributable to the hydroxyl proton is observed. In order to determine an equilibrium constant in the nmr, the spectrum of the pure monomer and the spectrum of the pure polymeric species must be known. Since in this work with 2-methyl-0-quinolinol and 8-quinolinol this information is not available, only qualitative information can be obtained. For this reason the following qualitative explanation of the data obtained from the infrared and nmr is given. The proposed equilibrium in equation (2) would result in a band in the infrared spectrum due to the monomer, around 3400 cm"1 (pyridine hydrochloride has an N+ H stretch at 3400-3500 cm 1). The presence of the dimer would result in a weakly associated hydrogen bonded species which could be detected by nuclear magnetic resonance spectroscopy. This model would be consistent with the nmr data, since at lower concentrations when only the monomer is present, the proton would not show a resonance signal in the vicinity of 9<f > at higher concentrations, however, the hydroxyl resonance in the dimer occurs around 8 o .
V
O Q H 00I I ©0 H (2)H
© IH 0 0
monomer dimer
42
In future work tenrperature should be used as a variable, instead of concentration, to study the equilibrium between the associated and non--associated states. Low temperatures will provide more information about the rate, of exchange between the associated and non-associated species. Also, other frequencies should be examined to see if indeed the monomer exists primarily as the proposed zwitter ion form.
!
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