• TABLE OF CONTENTS
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 Title Page
 Table of Contents
 List of Tables
 List of Figures
 Introduction
 Experimental
 Experimental data
 Discussion
 Summary
 Bibliography
 Acknowledgement
 Biographical note
 Copyright






Group Title: study of the infrared spectra of some substituted acetamides.
Title: A Study of the infrared spectra of some substituted acetamides.
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Title: A Study of the infrared spectra of some substituted acetamides.
Series Title: study of the infrared spectra of some substituted acetamides.
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Language: English
Creator: Letaw, Harry Jr.
Publisher: Harry Letaw, Jr.
Place of Publication: Gainesville, Fla.
Publication Date: June, 1952
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General Note: A dissertation presented to the Graduate Council of the University of Florida in partial fulfillment of the requirements for the degree of Doctor of Philosophy
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Table of Contents
    Title Page
        Page i
    Table of Contents
        Page ii
    List of Tables
        Page iii
    List of Figures
        Page iv
        Page v
    Introduction
        Page 1
        Page 2
        Page 3
        Page 4
        Page 5
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        Page 8
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        Page 13
        Page 14
        Page 15
        Page 16
        Page 17
        Page 18
        Page 19
    Experimental
        Page 20
        Page 21
        Page 22
        Page 23
    Experimental data
        Page 24
        Page 25
        Page 26
        Page 27
        Page 28
        Page 29
        Page 30
        Page 31
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        Page 57
        Page 58
        Page 59
        Page 60
    Discussion
        Page 61
        Page 62
        Page 63
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        Page 65
        Page 66
        Page 67
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        Page 92
    Summary
        Page 93
        Page 94
    Bibliography
        Page 95
        Page 96
        Page 97
        Page 98
        Page 99
    Acknowledgement
        Page 100
    Biographical note
        Page 101
        Page 102
    Copyright
        Copyright
Full Text












A STUDY OF THE INFRARED SPECTRA OF

SOME SUBSTITUTED ACETAMIDES





By

HARRY LETAW, JR.


A DISSERTATION PRESENTED TO THE GRADUATE COUNCIL OF
THE UNIVERSITY OF FLORIDA
IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE
DEGREE OF DOCTOR OF PHILOSOPHY








UNIVERSITY OF FLORIDA
JUNE, 1952














TABLE OF CONTENTS

Page

LIST OF TABLES ........................... it

LIST OF FIGURES ....... .. .. ... ... .. .. ili

Chapter

I INTRODUCTION...................... 1

II EXPERIMENTAL....... ..... .... ..... 20

III EXPERIMENTAL DATA ....,........... 4

IV DISCUSSION ................ ...... 61

V SUMMARY ........................ 93

BIBLIOGRAPHY ........................... 95

ACKNOWLEDGMENTS ...................... 100

BIOGRAPHICAL NOTE ............1......... 101














LIST OF TABLES


Table Page

I Abbreviations of Compounds . ........ 25

II Three Micron Region, N-dibutyl Series . 26

III Three Micron Region, N-butyl Series ...... 27

IV 5.7 Micron Band, N-dibutyl Series ........ 29

V The Six Micron Region. ..... ........ 30

VI 6.4 Micron Region, N-butyl Series . . . 31

VII Fluorine Stretching Bands . . . .. 33

VIII A Band and B Band Dilution Shifts, N-butyl Series 65












LIST OF FIGURES


Figure

1. Noise Level and CC14 Compensation Pattern ....

Z. N-butyl, Pure . . . . . . . . . .

3. N-butyl, 0.004M ....... .. ........ ..

4. N-butylChloro, Pure .................,


N-butyl

N-butyl

N-butyl

N-butyl

N-butyl

N-butyl

N-butyl

N-butyl

N-butyl


Page

36

37

38


Chloro, 0.006 M ..... ..........

Dichloro, Pure .... .. ..........

Dichloro, 0.006 M . . . . ... .

Chlorofluoro, Pure . ... .... .. .

Chlorofluoro# 0.005 M. ...........

Difluoro, Pure . . . . . .

Difluoro, 0.009M ..... .... .

Trifluoro, Pure . . .. . . . .

Trifluoro, 0.004 M .... . ...


5.

6.

7.

8.

9.

10.

11.

12.

13.

14.

15.

16.

17.

18.

19.


N-dibutyl, Pure&
N-dibutyl Pure . . . . . . . . . .



N-dibutyl Chloro, P.u .. . .. .. u. .

N-dibutyl Chloro. 0.007 ....... ......

N-dibutyl Dichloro, Pure
N-dibutyl Dichloro, Pure . . . . . .

N-dibutyl Dichloro, 0.008 M . ...M .,...

til


*


D

D


P











tv

Figure Page

20. N-dibutyl Chlorofluoro, Pure .... .. 5

ZI. N-dflutyl Chlo2rofluoro. O 0. 007.. M 56

22-0 Xdibuty1E Iiluero. Pure- i #..... 57

23. X-dibutyl Difluoro, 0. 007 ...... 0.. 0 8$



25. N-dibutyl Trifluoro. O.o0 6 ip .o.0........0 60











CHAPTER I

INTRODUCTION


Infrared spectrophotometry has taken its place among the more

useful tools at the disposal of the investigator studying inter- and intra-

molecular phenomena. This vast field has been thoroughly treated by

G. Herzberg in his two volume work "Molecular Spectra and Molecular

Structure", 3 Barnes and Bonner have briefly reviewed the history and

origin of infrared spectra in a reasonably elementary fashion.

The development of this technique has taken place from two points

of view which have only recently been fused in the literature. The em-

pirical approach is typified in a recent book, "Infrared Determination

of Organic Structures". 58 Theoretical aspects, on the other hand, are

discussed by Dennison in two papers22 23 which have appeared since

the discovery of the quantum mechanics. Barnes, Gorem Liddel and

Williams have assembled a bibliography of the literature of infrared

spectroscopy through the year 1943.9 In cooperation with the American

Petroleum Institute, the National Bureau of Standards is reproducing

reference spectra of a variety of compounds. Spectrograms are

periodically added to this publication

Because of the nature of the compounds dealt with in the present

study, much emphasis must be placed upon the concept of the hydrogen

I












bond or, as it is often termed, the hydrogen bridge. These two phases

will be considered to be interchangeable.

The phenomena which led to the postulation of the existence of

the hydrogen bond were observed primarily during cryoscopic and

ebullioscopic measurements. It was found that anomalous, or unex-

pected results were frequently obtained in the laboratory. Oddly

enough, the first postulation of the hydrogen bond was by Moore and

Winmill54 who, in 1912, found the concept necessary to explain the

weakness of the basic properties of trimethylammronium hydroxide.

Latimer and Rodebush39 are usually credited with recognizing

the generality of the hydrogen bond. In their development, however,

they erred in that they considered the hydrogen to be dicovalently

bound to its neighbors. Pauling56 found this to be incorrect after the

introduction of the quantum mechanics and placed the hydrogen bond on

a largely electrostatic basis. Though an occasional argument to the

contrary appears in the literature, it has become obvious that Pauling's

concept is basically correct.

Three rather thorough reviews on the general subject of hydrogen

bonding have appeared in fairly recent times. 19, 38, 60 The definitive

work on the subject is Pauling's "Nature of the Chemical Bond".55

Care must be taken in the study of the latter reference since the revision

published in the year 1948 is hardly complete up to that date.











3

Discussions of the amide group and its intermolecular linkages

have been centered largely about the problem of tautomerism in these

compounds. The origin of the proposal for a keto-enol tautomeric

shift in the amides is deeply bound in the macrochemistry of the group.

It has been found that the reactions of the amides occur in two different

categories. First, attack is possible upon the amide nitrogen itself.

This is typified by the reaction of the amides with rather strong bases.

One of the amine hydrogens can be directly replaced by a sodium atom.

This sodium salt may then be reacted with an alkyl iodide to yield the

N-substituted amide. Intermediately, the silver ion can react as the

sodium ion did or else demonstrate a second mode of attack by adding

to the carbonyl oxygen. When treated with an alkyl iodide, these com-

pounds react in such a way as to produce the N-substituted amide and

the imido ester, respectively. The extreme case appears on the treat-

ment of the free acid amides with dimethyl sulfate. Matsui53 reports

that the alkylation proceeds at temperatures below 1000 C. with the

recovery of the methyl hydrogen sulfates of the imido esters. The

present author believes that this reaction is not directly comparable

to the two previously described in that its mechanism is undoubtedly

quite different.

The reaction schemes above have demonstrated essentially the

liability of the proton attached to the amide nitrogen. This fact is










4

startling in view of the apparent availability of the "electron pair" on

the electronegative nitrogen. Comparison of the amides with the

amines led to the rather ill-considered opinion that the former com-

pounds should be basic. The amides are actually feebly acidic except

in the most extremely acid solvents. It is this apparent contradiction

which has led to many studies of the amide linkage.

It became evident that purely chemical methods would hardly suf-

fice for the final resolution of the problem of the structure of the

amide group. Recognition of the fact that this linkage is intimately

related to the overall question of the nature of the proteins and poly-

amides served as a tremendous stimulus to the development and

improvement of means of attack upon this problem.

Kumler and Porter37 applied data obtained from dipole moment

measurements to the problem of the molecular configuration of acetam-

ide, N-ethyl acetamide, N-diethyl acetamide and N-dimethyl acetamide.

Utilizing the method and apparatus of Williams and \7eissberger,72

they calculated the proportion of the classical enolic tautomer present.

It was found that this fraction is vanishingly small even for acetamide

itself. They stated that the true configuration, at least. in benzene solu-

tion, is a resonance hybrid of the classical ketonic amide and an approx-

imately ten percent contribution from an excited form. This excited

form was drawn in such way that the nitrogen is doubly bonded to the










5

carbon and the oxygen is singly bonded to the carbon. The nitrogen and

oxygen possess formal positive and negative charges, respectively.

The first application of the infrared spectrophotometer in this
26
problem was by M. and R. Freymann in 1936. These researchers

noted that in a keto-enol type tautomeric mixture, the characteristic

hydroxyl valence bond stretch would appear in the absorption spectrum

and that its intensity would correspond to the proportion of enolic form

in the mixture. Since they were unable to find this band in the spectra

of the liquids, melts or vapor phases of several amides, they concluded

that the classical tautomer does not exist. The Freymanns postulated

an intramolecular association. Their final picture resulted from ascrib-

ing a partial double bond character to the carbonyl and carbon-nitrogen

bonds and attaching a partial no-bond character to the hydrogen bridge

between nitrogen and oxygen. The implication of their structure is that

the two sets of bonds are approximately symmetrized with respect to the

bond strengths. This conclusion is amplified slightly and reiterated by

R. Freymann27 and by M. Freymann in publications appearing very

early in 1939.

Raman spectral data have also been of use to workers studying

problems involving the amides. The general background material is

covered adequately by Herzberg. 32 In 1937, Ananthakrishnan3 inter-

preted the Raman spectra of acetamide and propionamide as indicating










6

the existence of the classical structure plus an excited state differing

from the classical tautomer in that the proton was not actually trans-

ferred from the nitrogen. This second structure was an approach to

that of the Freymann's26 mentioned above; however, in it was implied

a somewhat more sophisticated grasp of the pattern involved in reso-

nating structures.

During this time, Buswell, Rodebush and co-workers had been

studying molecular association phenomena by means of the infrared

spectrometer. In 1938, their first attack on the amide problem

appeared.14 It was concluded that the simple amides are associated

in indeterminately large polymers, the structure of which they were

unable to deduce. The N-disubstituted amides showed no evidence of

association. Because of the spectroscopically obvious perturbations

of the nitrogen-hydrogen bond in the N-monosubstituted amides, it was

deduced, following the suggestion of Copley, that these structures form

a cyclic dimer through two hydrogen bridges. Feeling the necessity of

showing the existence of some tautomeric form, these workers stated

that in the cyclic dimer there is an effective transfer of the protons to

the carbonyl oxygen with the simultaneous assumption by the carbon-

nitrogen linh of double bond character.

Sannie and Poremski61, apparently in ignorance of the work of

Buswell, Rodebush and Roy14 and that of Ananthakrishnan3, deduced










7

approximately the same structures for the amide linkage as did these

authors. They found that the amides associate to a great extent in

benzene solutions, but very little in water and in ethanol. For the polar

solvents, they extended the idea of M. and R. Freymann 26 concerning

an internally symmetrized form of the hydrogen bridge, by applying

independently the idea of Ananthakrishnan3 and justifying it with a

rather unsatisfactory steric argument. Sannie and Poremski believed

that the amides and N-monosubstituted amides in the non-polar solvent

form dimeric or trimeric rings through hydrogen bonding. In the case

of the N-disubstituted amides, an internal ring involving hydrogen bonding

of an alkyl hydrogen to the carbonyl oxygen was suggested. It is inter-

esting to note that each form presented by Sannie and Poremski was

essentially a resonance hybrid of the tautomeric form and a quasi-tauto-

meric form in which the nitrogen retains the hydrogen and a positive

charge, while the oxygen is singly bound and negatively charged.

The studies by Buswell's group continued for several years in the

more general aspects of association. During the year 1940, several

significant advances arose from the work of this group. Buswell,

Downing and Rodebush12 stated that they had eliminated the problem of

the tautomeric shift previously postulated for the dimeric form of the

N-monosubstituted amides. They observed that the characteristic

absorption of the carbonyl group in high dilutions in carbon tetrachloride










8

is essentially the same as that observed in the spectra of the pure

liquids. This led to the realization that, in the dimer, there is the pos-

sibility of cis-trans isomerization with respect to the orientation of the

aryl or alkyl substituents on the a mide nitrogen. They then proposed

that only the trans form of the dimer is stable, presumably for steric

reasons. Introducing a note of indeterminism, they stated that the

characteristic frequencies of the nitrogen-hydrogen valence stretch in

which the hydrogen was bridged either to another nitrogen or to an

o::ygen are identical.

Busvwell, Krebs and Rodebush, 13 on the basis of the preceding

work, stated that the modes of association in the proteins are through

amino hydrogen bridges to oxygens. This paper also contained the

heuristic idea that the characteristic perturbed bands of the amino

hydrogen link occurring at 3. 00 microns (3333 cm"1) and 3.22 microns

(3125 cm-1) are due to linear and cyclic dimers# respectively. It was

suggested that the difference in frequencies associated with these forms

lies either in resonance stabilization of the ring or in the cis-trans iso-

merization mentioned above.

Lecompte, collaborating with R. Freymann 40s41 extended infra-

red studies of the amides into the six to nineteen micron (1667-526cm 1)

range of the spectrum. Buswell, Downing and Rodebush12 had utilized

an isolated datum from that region, but had neglected the golden












opportunity which was at hand by continuing to concentrate upon the

three micron region. In the first paper by Lecompte and Freymann,

there is a fairly complete discussion of the theories of the amides

which had been advanced prior to that time. Once again, a case is

encountered in which the authors were unfamiliar with the literature.
The paper by Busvwell, Krebs and Rodebush 3 are unknown to them.

They reiterated that the excited dimer of Sannie and Poremskit1 is an

improvement upon the form proposed by Buswell, Rodebush and Roy;14

however, they rather vehemently denied the validity of the inner tau-

tomer of the former. Lecompte and Freymann foundthe cyclic tauto-
mer of Buswell, Rodebush and Roy14 to be satisfactory, being totally
unaware of the fact that this form had been discarded in a later pub-

lication. 1

In an X-ray diffraction study of the acetamide molecule# Senti'

and Harker62 found that the molecule is planar. This result is predic-

ted by either the ketonic or enolic form. It was found that the carbon-

oxygen bond distance is 1. 28 A., the carbon-methyl distance, 1. 51 A,
and the carbon-nitrogen distance, 1.38 A. The first measurement
agrees well with lata from formic acid, oxalic acid dehydrate, dilreto-
piperazine and glycine, while the second value is normal. The carbon-
nitrogen bond distance is somewhat greater than the distance reported

for diketopiperazine, but is approximately the same as that obtained

for urea and thiourea. The oxygen-carbon-methyl angle is 1290, the












nitrogen-carbon-methyl, 1090, and the nitrogen-carbon-oxygen, 122O.

It was found that a nitrogen-hydrogen-oxygen bridge exists in the mole-

cule. Its length is 2.86 A. Senti and Harker conclude that the nitro-

gen-hydrogen bonds are in the plane of the molecule and that only

the ketonic form exists in the crystal.

Little more than confirmatory work was done during the remain-
ing war years. A tremendous quantity of data pertaining to the amide

problem was collected in connection with the effort to synthesize peni-

cillin. The germane portion of this work is reported in a previously

cited reference. 58

Because of the immense productivity of Henri Lenormant, his

publications i4Z-51 will be described from his excellent review paper.

i'Infrared Spectra of the Peptide Linkage".51 In order to utilize this

review to its maximum extent, it is necessary to step out of chronol-

ogy for the examination of two papers appearing during this interval.

Richards and Thompson59 stated that in their opinion, the N-

monosubstituted amides could exist in two mesomeric forms com-

posed of four resonating structures, the classical amide and its

tautomer and the corresponding excited states. From these deductions,

Richards and Thompson felt that the characteristic amide vibrations

should be the valence stretches and deformations of the carbon-ox-ygen

double bond, carbon-nitrogcn double bond, hydrogen-oxygen bcnd and

hydrogen-nitrogen bond.











11

In their study, emphasis is placed for the most part on the

absorption occurring in the five to seven micron (2000-1429 cm-1)

range; however, they did study the three micron region briefly. They

concluded that bands occurring in the latter range are actually caused

by various nitrogen-hydrogen stretching modes, but that the hydroxyl

stretch is not observed. This, of course, eliminated the possibility

of the classical enolic form and its resonance hybrid. Other bands

in the three micron region were attributed to hydrogen bridges be-

tween two atoms of nitrogen and between nitrogen and oxygen.

The most interesting phenomena reported were in the lower fre-

quency region. Two bands were observed in this spectral region of

the amides and the N-monosubstituted amides whether they were in the

solids fused or liquid state or in solution. The one appearing at the

higher frequency, about 1667 cm'" (6.0 microns), was denoted "A"

and that at 156Z cm"1 (6.4 microns), "B". These have been labeled

the "Amide I" and "Amide II" bands, respectively, by other authors.58

It was found that the A band occurs in all amides regardless of

the state of substitution upon the nitrogen. This band is displaced to

slightly lower frequencies when compared to the carbonyl band of other

substances, but it is sufficiently close to that absorption to be classi-

fied. The B band does not occur in the N-disubstituted amides. It is

shifted to rather lower frequencies in the case of the N-monosubstituted










12

amides than those observed in the non-substituted amides. It was

found that upon dilution, the A band shifts to higher frequencies and the

B band to lower. Substitution of an electrophilic group upon the nitro-

gen tends to reinforce the A band and weaken the B band. This is

further evidence that the A band can be identified as the carbonyl

stretch.

They then proposed four possibilities for the origin of the B band.

First, the ketonic carbon-nitrogen bond with some double bond charac-

ter; second, the enolic carbon-nitrogen bond with somewhat less than

one-hundred percent double bond character; third, the amino hydrogen

deformation; fourth, an overtone or combination band. They argued that

if this band is assigned to the enolic form of the molecule, it is tacitly

assumed that both tautomers exist in the solid or pure state and in solu-

tions. This is contrary to the evidence presented by Senti and Harker.

That it is the somewhat doubly bound carbon and nitrogen in the ketonic

form was considered to be unlikely since this form would be expected to

exist to a higher degree in the N-disubstituted compounds than in the

N-monosubstituted. The B band is not observed in the former class of

compounds. The shape and intensity of the band tend to rule out the

possibility of its being an overtone or combination band. Thus it was

presumed that this band is caused by the nitrogen-hydrogen deformation.










13

Hartwell, Richards and Thompson, 30 in a general investigation

of the characteristic absorption frequency of the carbonyl group,

reported that the N-disubstituted aides absorb in the neighborhood

of 1710 cm* (5. 85 microns). In studies of the chloroacetic acids, they

observed that the substitution of such electrophilic groups shifted the

absorption of the carbonyl group to higher frequencies.

Returning now to the work of Lenormant, 5 it is found that in

previous investigations, he had observed the so-called B band in the

amides, but had not studied its sensitivity to dilution. In the course of

his researches, he came to the conclusion that the B band results from

a symmetrization of the carbon-nitrogen and carbon-oxygen bonds in

that the carbonyl bond has somewhat less than one-hundred percent

double bond character and the carbon-nitrogen link approaches that un-

known percentage double bond character. This conclusion is reinforced

to a great extent by the fact that only the A band occurs in N-.bromo-

acetamide, only the B band in sodium acetamide and both bands with

equal intensities appear in N-ethyl acetamide. It was noted earlier

that the silver salt of acetamide plays a dual role in substitution reac-

tions. Lenormant found that the substitution of silver upon the nitrogen

in acetamide yields a strong B band and a relatively weak A band which

is displaced to somewhat lower frequencies.










14

It was found that in the spectra of the lactams, cyclopeptides and

diketopiperazine, the B band does not appear. Observing that the

nitrogen-hydrogen deformation should be found in each of these com-

pounds, Lenormant felt that the explanation of the origin of this band

by Richards and Thompson is invalid. On basic hydrolysis, the cyclo-

peptides lose the A band and gain the B band, while diketopiperazine

presents a very intense B band and a weakened and shifted A.band. In

the polyamides, such as the nylons, the A and B bands both occur as

long as there is an amino hydrogen present. Replacement of this atom

by an alkyl group results in the disappearance of the B band.

On deuteration of the compounds possessing the B band, it is

found that the band shifts to lower frequencies by a ratio of 1/1. 05. In

other less ambiguous cases, replacement of a hydrogen atom by a deuter-

lum atom results in shifts of the frequency very nearly in a ratio of

1/1. 33. 3- The implication here is that the B band is not strictly a

function of the nitrogen-hydrogen defurmation as was stated by Richards

and Thompson.59

In view of the fact that only the A band appears in the N-disubsti-

tuted amides and that this band is shifted to lower frequencies with

respect to the normal absorption of the carbonyl group, Lenormant

concludes that those compounds exist in a mesomeric state to

which the excited form makes a contribution of approximately











15

fifteen percent. He arrives at a similar conclusion for the non-

substituted amides.

With respect to the N-monosubstituted amides and the peptides,

he concludes that the B band exists only if the group substituted upon

the nitrogen can assume a position cis to the oxygen. This is exem-

plified by the N-monosubstituted amides, peptides and polyarnides.

This band fails to appear if this orientation cannot be attained as is

the case for the lactams and diketopiperazine.

Lenormant mentions the beta-di-ketoncs and beta-diketonic

esters. When the sodium, copper or manganous salt of any of these

compounds is formed, the carbonyl stretch shifts to markedly lower

frequencies. Sodium tends to shift the band to the greatest extent.

This is compared to the phenomena observed following the substitution

of bromine, sodium or an alkyl group for the amino hydrogen in

N-monosubstituted aides and is found to support his idea of the con-

figuration of the amide group.

Lenormant clarified the situation in the three micron region

to a great extent by obtaining the infrared spectra of N-deuterated

amides. He found that all of the bands in that region previously re-

ported to be nitrogen-hydrogen stretching modes shifted the proper ratio.

This, of course, could not serve as a basis for the elimination of

the tautomeric form. He agreed with the work of Buswell, Krebs











16

and Rodebush13 in the assignment of bands in this region to linear

and cyclic di-ners.

Cromwell, EMiller, Johnson, Frank and V-aIllpce prepared a

series of amino-substituted, alpha-beta unsaturated ketones. Numer -

ous anomaliess" appeared in the infrared spectra of these compounds.

Because of the complexity of these compounds, it is difficult to assign

many of the bands appearing. It is significant to note that a strong or

very strong band occurs in the region in which the previously de-

scribed B band appears. The characteristic carbonyl absorption is

shifted to lower frequencies than those observed in the simple dialkyl

ketones. The B band is found in the two cases in which the amino

group is substituted in the alpha position; however, the nitrogen is

either part of a morpholino group or else it is methyl. benzyl substi-

tuted. The carbonyl group appears at about the same frequency at

which it is found in the N-disubstituted amides. In the beta amino

compounds, the B band appears regardless of the state of substitution

of the nitrogen. The A band occurs at much lower frequencies than

would be expected in the corresponding amides. The authors were

unable to draw any definite conclusions regarding the nature of these

bands.

Darmon and Sutherlandl8 found that the presence of water in pro-

teins used in infrared spectroscopic work does not affect the nitrogen-











17

hydrogen stretching frequencies. They confirmed the previously

reported supposition13' 51 that the 3060 cm"I (3.27 micron) band is

characteristic of the cyclically dimeric form of the amides and that

the 3280 cm-1 (3.05 micron) band is characteristic of the linear dimer.

In the year 1951, the Discussions of the Faraday SocietyZ4 con-

tains a series of interesting comments regarding the general problem

of the amides. On page 274. Sutherland stated that the characteristic

absorptions of the nitrogen-hydrogen stretches perturbed by the car-

bonyl group are in the range 3320-3240 cm-1 (3.01-3.09 microns).

When bonded to another nitrogen, the absorption occurs from 3350-

3150 cm"1 (3.03-3.17 microns). Lenormant suggested on page 319

that the two bands appearing in the three micron range and the A and

B bands might be caused by two distinct forms of the amides.

In a thorough study of acetamide, Davies and Hallam0 found

that this compound exists almost wholly as a cyclic trimer in chloro-

form solution, a dimer in acetone and a monomer in methyl cyanide.

In the trimer, by virtue of the two distinct vibrational modes shown

for the amino hydrogen stretch, it is assumed that only one of the

hydrogens of each amino group is involved in ring formation. They

found the carbonyl stretch to be at 1700 cm"1 (5. 88 microns) in the

monomeric form and at 1678 cm-1 (5.96 microns) in the associated

form. A very strong band occurring at 1595 cm"1 (6.27 microns)











18

is assigned to the deformation of the unassociated amino group follow-

ing Richards and Thompson.59

Bates and Hobbs10 have recently investigated the dipole moments

and group structure of some acid amides. In an effort to explain the

nature of the group without utilizing the concept of resonances they

assumed the existence of a planar structure and proceeded to discuss

the observed moments on the basis of dipole interactions* Their deduc

tions are somewhat vitiated by comments by Kumler;36 however, the

conclusion that the amides exist in the ketonic form is in good agree-

ment with the numerous investigations presented heretofore.

A reasonably. consistent interpretation of the work which has been

presented above would lead one to believe that the amides exist in an

associated form in the solid state& melt, pure liquids and fairly con-

centrated solution in non-polar solvents. They are neither dimerized

in polar solvents nor associated in dilute solutions in non-polar solvents.

Further, the amide group must be thought of as being planar and ketonic

in form.; If one assumes that the enolic form may exist, it must be

only as a result of the symmetrization of a hydrogen bridge between the

amino hydrogen and the carbonyl oxygen. It is assumed that hydrogen

bridging to nitrogen does not tend to occur. The Bland is found in the

spectra of certain amides. This band has not been explained on the

basis, of the present information.











19

Previous work in this field has concentrated upon the variation

of the substituents upon the nitrogen. One finds references made to

benzamide and, occasionally, to amides of the higher aliphatic acids;

however, little work has been done using amides in which the hydrogen

of methyl or methylene groups alpha to the carbonyl group have been

systematically replaced by electrophilic groups. It is evident that this

substitution would inhibit the formation of An essentially single-bonded

configuration of the carbonyl group. It would encourage the deformation

of the electronic cloud about the nitrogen in such a way as to promote

the assumption of partial double bond character by the carbon-nitrogen

bond. Thus, it should be possible, these shifts being predictable, to

create an amide with sufficient abnormality that the questionable absorp-

tions of this group could be identified.

This study was carried out on a series of N-butyl-ethanamides and

one of NN-di-butyl ethanarmides. The alpha-methyl group in each series

is step-wvise halokcnatcd. The compounds investigated in this research

are listed in Table I.











CHAPTER II

EXPERIMENTAL

Two infrared spectrophotometers were utilized in this research.

The first was a Perkin-Elmer Model 12-C single-beam instrument in

the Naval Stores Research Laboratories of the Glidden Corporation in

Jacksonville, Florida. A Perkin-Elmer Model 21 double-beam infra-

red spectrophotometer in this laboratory was used for the final phases

of this research. The frequencies of all bands are reported as

obtained from data gathered by use of the latter instrument because

of its superior calibration characteristics.

The electrical and optical details of both instruments are com-

pletely described in manuals and drawings prepared by the Perkin-

Elmer Corporation. Further data concerning the double-beam instru-

ment are available in a series of papers by White and Listron.69-71

The method of calibration of the instruments is mentioned in the

previously cited manuals. A table of wavelengths and wavelength

standards suitable for such a calibration was published by Plyler and

Peters.57

Experimental techniques are described in a number of sources.

Two previously mentioned books958 contain valuable information on

this subject. A rather general paper concerning both instrumentation

20











21
73
and techniques has been published by Williams. Specific applications

of the double-beam instrument are discussed in a later paper.

The :methods used in the present study are essentially conformist.

It was found that concentrations could be rapidly determined by prepa-

ration of several solutions of known concentration and obtaining their

spectra. Comparison of the carbonyl band absorption intensities of

the known solutions to those of the roughly prepared solutions of the

same order of magnitude of dilution permitted this determination.

The compounds used were obtained from two sources. All of

the acetamides en:cept for the trifluoro derivatives were prepared by

Fried28 under the direction of Tarrant in this laboratory. The latter

compounds were prepared by Tarrant and Letaw. 66 Reagent grade

carbon tetrachloride was used as the solvent in all dilution work. In

each case, the absence of extraneous bands in the infrared spectro-

gram was the criterion of purity.

The spectra used in this research were all obtained under iden-

tical conditions with respect to instrumentation. Scanning was at the

rate of one micron per minute. Automatic slit program IV was used.

This provided slit widths of 14, 37 and 111 microns at wavelengths of

1. 8, 5 and 10 microns (5556, 2000 and 1000 emr"l), respectively. In

the original work, the scale of recording was two inches per micron;











22

however, for the purpose of reproduction, this was reduced to one

inch per micron.

The spectral region studied was from two to fifteen microns.

These limits were imposed by the sodium chloride optics used. Work

was carried out at 250 C. in a relative humidity of less than fifty

percent. As a result of compensation for the solvent, carbon tetra-

chloride, in dilution work, the sensitivity of the instrument was

reduced considerably in the region beyond ten microns. For this

reason, discussion of the spectra studied will be limited to the 2-10

micron (5000-1000 cm1) interval.

Spectra were obtained for the pure materials as well as for

several dilutions of each compound. In general, further dilutions were

not prepared if there were no change in the spectrum for two consec-

utive dilutions. Because it is a solid at 250 C., N-butyl-2, 2-dichloro-

ethanamide was not run as the pure substance. A very small quantity

of carbon tetrachlorTde was added to it in order that a homogeneous,

non-scattering mull might be obtained.

For the pure compounds, a demountable cell consisting of two

salt plates pressed together over the liquid was used. The thickness

of the sample was adjusted by variation of the pressure upon the plates

in order to obtain a spectrograph of suitable intensity. Dilutions were

carried out in cells approximately 0.025, 0. 100 and 0.500 mm. thick.











23

All dilutions reproduced in this paper were obtained in a cell 517.8

microns thick compensated with a cell containing a layer of carbon

tetrachloride 486.6 microns in thickness. The degree of compensa-

tion obtained is shown in Figure I.











CHAPTER HI

EXPERIMENTAL DATA

Reproductions of the spectrograms obtained in this research are

attached at the end of this chapter. Only the spectrum of the pure

compound and one dilution are reproduced. This concentration,

approximately 0.01 M, is of such a magnitude that no further shift

in the frequency of any band is observed upon further dilution. The

value of each dilution and the name of the compound is indicated on

each of the figures. The bands will be discussed in order of decreas-

ing frequency or, equivalently, increasing wavelength.

Because of the rather lengthy names of the compounds investi-

gated, abbreviations have been constructed for them. There is little

possibility of ambiguity in the abbreviated nomenclature as it appears

in Table I.

To facilitate the description of the bands in the spectra at hand,

the following notation will be utilized to indicate relative band inten-

sities: vs e very strong, s = strong, m medium, w = weak, and

vw a very weak.

In all of the compounds investigated, a broad, weak absorption

band was found in the 2.25-2.75 micron (4444-3636 cm" ) region. It

can be ascribed either to an overtone or combination band or to a













small amount of water existing as an impurity in the compounds.

Because of its contour, the former explanation is the more probable;

moreover, it will be recalled that Darnmon and Sutherland have estab-

lished the fact that water does not affect the characteristic amide

absorption bands.

TABLE I

ABBREVIATIONS OF COMPOUNDS

Name of Compound Abbreviation

N-butyl-ethanamide N-butyl

N-butyl-2-chloro-ethanamide N-butyl chloro

N-butyl-2, 2-dichloro-ethanamide N-butyl dichloro

N-butyl-2-chloro-2-fluoro-ethanamide N-butyl chlorofluoro

N-butyl-2 2-difluoro-ethanamide N-butyl difluoro

N-butyl-2, 2# 2-trifluoro-ethanamide N-butyl trifluoro

N, N-di-butyl-ethanamide N-dibutyl

NO N-di-butyl-2-chloro-ethanamide N-dibutyl chloro

N, N-di-butyl-2, 2 -dichloro-ethanamide N-dibutyl dichloro

N, N-di-butyl-2-chloro-Z -fluoro-ethanamide N-dibutyl chlorofluoro

N. N-di-butyl-2, 2-diflporo-ethanamide N-dibutyl difluoro

N, N-di-butyl-2, 2, Z-trifluoro-ethanamide N-dibutyl trifluoro











26

The N-dibutyl compounds present a rather puzzling set of ab-

sorptions in the region immediately around three microns. These

bands and their relative intensities are noted in Table II. On dilution,

there is no evidence of shifts in these bands. The bands appear to be

overtones or combinations. With respect to impurities, the corres-

ponding N-butyl compounds could be present; however, the rather

large differences in boiling points relative to the N-dibutyl compounds
28 66
makes this possibility rather unlikely.Z8

TABLE II

THREE MICRON REGION, N-DIBUTYL SERIES

Compound Wavelength Frequency
(microns) (cm1-)

N-dibutyl 2.91 (w) 3436
3.07 (vw) 3257

N-dibutyl chloro 2.90 (w) 3448
3.04 (w) 3287

N-dibutyl dichloro 2.95 (vw) 3390
3.04 (w) 3289

N-dibutyl chlorofluoro 2.93 (vw) 3413
3.05 (w) 3279

N-dibutyl difluoro 2.88 (vw) 3472
3.03 (vw) 3300

N-dibutyl trifluoro 2.88 (vw) 3472
3.01 (w) 3322











27

The bands occurring in the neighborhood of 2.9 microns

(3448 cm-1) correspond in frequency to the unperturbed hydrogen-

nitrogen stretch found in the N-monosubstituted amides. Aside from

the fact that such impurities are unlikely, it must be recognized that

virtually none of the nitrogen-hydrogen links would exist in an un-

perturbed state under the concentration conditions existing in the

system.

TABLE In

THREE MICRON REGION. N-BUTYL SERIES


Compound


N-butyl


N-butyl chloro


N-butyl dichloro


N-butyl chlorofluoro


N-butyl difluoro



N-butyl trifluoro


Wavelength
(microns)
Pure Dilute
3.05 (s) 2.91 (w)
3.24 (m)

3.07 (s) 2. 93 (w)
3.26 (m)

3.07(s) 2.92 (w)
3.26 (m)

3.06(s) 2.92 (w)
3.26 (m)

2. 92 (m) 2.92 (w)
3.06 (s)
3.24 (m)

3.04 (s) 2.92 (w)
3.24 (m)


Frequency
(cm-1)
Pure Dilute
3279 3436
3086

3257 3413
3067

3257 3425
3067

3268 3425
3067

3425 3425
3268
3086

3289 3425
3086


Shift In
Frequency
(cm-1)
/ 157
S350

/ 156
74 346

/ 168
/ 358

/ 157
358

0
7/ 157
/ 339

4- 136
/- 339











28

As will be shown later, the carbonyl absorption frequencies of

the N-dibutyl series occur in the range 1650-1680 cm-1 (6. 06-5. 95

microns). The first overtone of this absorption should occur in the

3300-3360 cm-1 (3.03-2. 98 micron) region or at slightly lower fre-
35
quencles. It would seem logical to identify the bands occurring

around 3.00 microns (3300 cm-1) as the first overtone of the carbonyl

group. By the same reasoning, the much weaker absorptions at

slightly higher frequencies could be overtones of a band occurring at

only slightly shorter wavelengths than that of the carbonyl. This

band is described later.

The three micron region of the spectra of the N-butyl amides is

extremely rich in significant bands. The pure compounds all have

two well-defined bands which may be identified as perturbed nitrogen-

hydrogen stretches. In additions N-butyl difluoro presents a very

sharp absorption at 2i92 microns (3425 cm-1) This is doubtless due

to the unperturbed nitrogen-hydrogen stretch. As the solutions be -

come more dilute, the two perturbed bands are markedly weaklened

and the unperturbed band is strengthened, In every case, it is found

that the lower frequency band disappears before the higher one.

Observed frequencies and intensities are recorded in Table II.











29

In both series of compounds, the carbon-hydrogen stretch pre-

sents three absorptions in the neighborhood of three and one-half

microns. A shoulder of medium intensity occurs at 3.42 microns

(2924 cm-1), a strong band is found at 3.43 microns (2915 cm"1) and

a medium band appears at 3.51 microns (2849 cm-1). There is no

apparent change in the bands of the compounds into which halogen

substituents have been introduced. This is, of course, due to the

overwhelming intensities of the carbon-hydrogen valence stretches

attributable to the butyl group.

TABLE IV

5.7 :.MICRON BAND, N-DIBUTYL SERIES

Compound -Wavelength Fr e qucncy
(microns) (cm- )

N-dibutyl 5. 79 (vw) 1727

N-dibutyl chloro 5.77 (w) 1733

N-dibutyl dichloro 5.70 (w) 1754

N-dibutyl chlorofluoro 5.65 (w) 1770

N-dibutyl difluoro 5.67 (w) 1764

N-dibutyl trifluoro 5.63 (vw) 1776

Another unusual band appears In the spectrograms of the N-

dibutyl series at approximately 5.70 microns (1754 cm1-). These
bands have not been previously reported. The frequency and










30

intensity data appear in Table IV. It is interesting to note that all

of the bands are sharp except for those of the N-dibutyl chloro and

N-dibutyl trifluoro amides. There is no shift of frequency upon dilu-

tion. It is possible that these are overtone or combination bends.

TABLE V


THE SIX MICRON REGION

Compound Wavelength
(microns)
Pure Dilute

N-butyl 6.04 (vs) 5.93 (vs)

N-butyl chloro 6. 02 (vs) 5.93 (vs)

N-butyl dichloro 5.98 (vs) 5.86 (vs)

N-butyl chlorofluoro 5.96 (vs) 5.84 (vs)

N-butyl difluoro 5.94 (vs) 5.82 (vs)

N-butyl trifluoro 5.86 (vs) 5.76 (vs)

N-dibutyl 6.07 (vs) 6.07 (vs)

N-dibutyl chloro 6.03 (vs) 6.04 (vs)

N-dibutyl dichloro 6.00 (vs) 5.94 (s)
6.03 (vs)

N-dibutyl chlorofluoro 5.97 (vs) 5.93 (s)
6.01 (va)

N-dibutyl difluoro 5.97 (vs) 5.93 (m)
6.00 (vs)

N-dibutyl trifluoro 5.92 (vs) 5.91 (vs)


Frequency
(cm-1)
Pure Dilute

1656 1686

1661 1686

1672 1706

1678 1712

1684 1718

1706 1736

1647 1647

1658 1656

1667 1684
1658

1675 1686
1664

1675 1686
1667

1689 1692


Shift In
Frequency
(cm-1)











31

Table V contains the data obtained for the absorptions occur-

ring in the immediate neighborhood of six microns. Each of the

compounds presents an extremely strong absorption in this region.

It has become rather obvious in the course of previous work that

these bands are characteristic of the carbonyl group.

TABLE VI

6.4 MICRON REGION, N-BUTYL SERIES


Compound



N-butyl


N-butyl chloro


N-butyl dichloro


N-butyl chlorofluoro


N-butyl difluoro


N-butyl trifluoro


Wavelength
(microns)
Pure Dilute

6.41 (vs) 6.66 (ve)
6.45 (s)

6.40 (s) 6.60 (a)
6.47 (vs)

6.40 (s) 6.61 (s)
6.47 (s)
6.55 (vs)
6.40 (s) 6.60 (s)
6.47 (vs)

6. 39 (s) 6.57 (s)
6.44 (vs)

6.40 (vs) 6.57 (s)


Frequency
(cm-1)
Pure Dilute

1560 1502
1550

1562 1515
1546

1562 1513
1546
1527
1562 1515
1546

1565 15Z4
1553

1562 1524


Shift In
Frequency
(cm-1)

58
48

47
31

49
33
14
47
31

41
29

41


By far the most intriguing band to be found in the infrared spec-

tra of the amides and amide-like compounds is that occurring in the

neighborhood of 6.40 microns (1562 cm-1) in the spectra of the











3Z

N-monosubstituted amides. This band has been discussed at length

in Chapter I. Deductions concerning the data obtained in this research

will be presented in the following chapter. The data are listed in

Table VI.

A fairly strong absorption band was found in the neighborhood of

6,85-7. 10 microns (1460-1408 cm-1) for all of the compounds inves-

tigated. On the basis of theoretical band assignments for hydrocar-

bons, 32 this band is assumed to be caused by the asymmetric defor-

mations of the methyl and methylene groups. It is observed that the

band is much more intense in the N-dibutyl compounds than it is in

the N-butyl derivatives

It was found that a band occurred at about 71 25 microns (1379

cm"1) in all cases, This band has been identified by Davies and
20
Hallam6 specifically for acetamide6 as the carbon-nitrogen stretch.

In view of the fact that there is no change in the frequency of this

band upon diiutions this assignment is not justified in the present

case* In subsidiary spectrograms obtained in the course of this

research it was found that this band occurs in n-butanol and in the

three n-butyl, n-hexyl and n-octyl amines with virtually no change

in form or in relative intensity. It is believed that this band and

those occurring in the neighborhood of 7.50 microns (1333 cm*l)













are probably due to symmetric modes of deformation of the various

carbon-hydrogen configurations in these compounds.

TABLE VII

FLUORINE STRETCHING BANDS


Compound


Wavelength
(microns)
Pure Dilute


Frequency
(cm-1)
Pure Dilute


Shift In
Frequency
(cm-1)


N-butyl chlorofluoro




N-dibutyl chlorofluoro





N-butyl difluoro




N-dibutyl difluoro


N-butyl trifluoro



N-dibutyl trifluoro


8.98 (m)
9.20 (,)
9.42 (vs)
9.65 (m)

8.98 (m)
9.14 (vs)
9.30 (s)
9.47 (m)
9.77 (w)

8.80 (vs)
9.03 (vs)
9.19 (vs)
9.39 (vs)

8.95 (vs)
9.43 (vs)

8.28 (vs)
8.46 (vs)
8.60 (ve)

8.32 (vs)
8.48 (vs)
8.77 (va)
8.3 5 (vs)


9.26 (w)
9.51 (vs)


8.98 (w)
9.12 (m)
9.42 (vs)

9.77 (w)

9.01 (vs)
9.03 (vs)
9.22 (vs)
9.39 (vs)

9.04 (s)
9.47 (vs)

8.32 (s)
8.60 (vs)
8. 60 (vs)

8. 30 (vs)
8.48 (s)
8.77 (vs)
8. 8* (s)


1112
1087
1062
1036

1112
1094
1075
1056
1024

1136
1107
1088
1065

1117
1060

1208
1182
1163

1202
1179
1140
1130


1080
1052


1112
1096
1062

1024

1110
1107
1085
1065

1106
1056

1202
1163
1163

1205
1179
1140
1126


- 10
- ---

0
S2
- 13

0

- 26
0
- 3
0

- 11
- 4

- 6
- 19
0

,3
0
0
- 4











34

A band found in the 8.00-8.20 micron (1250-1Z19 cm"1) range

is presumed to be attributable to rocking or wagging modes in the

carbon-hydrogen links. In some of the fluorinated compounds, this

band is partially masked by carbon-fluorine valence stretches.

In view of the pronounced hydrogen bonding possibilities

through fluorine atoms, it is not surprising that dilution produces

considerable changes in the absorption frequencies of the carbon-

fluorine bonds. In Table VII, the data for the absorption attribu-

table to the valence stretch of the carbon-fluorine bond are given.

The carbon-fluorine stretch assignments are in the 8.28-9. 77

micron (1208-1024 cm-1) range. It is evident that several of the

non-fluorinated. compounds absorb radiation in that band of frequen-

cies. This leads to some ambiguity in the selection of the bands;

however, it is seen that the bands cited in Table VII are, for the

most part, vastly more intense than are those found in this region

of the spectra of the other six compounds. The intensity of an infra-

red band is, to a very good first approximation, proportional to the

change of the dipole moment produced by the vibrational mode which

is excited. The stretching modes of the carbon-fluorine vibrations

are found in this region. It can be said that the intensities of these

bands should be! much greater than those of the less active deforma-











35

tional modes of the carbon-hydrogen bonds which also occur here.

Thus, it is reasonable to assume that these extremely intense bands

are due to carbon-fluorine valence stretches.














MICRONS
35 .5 6 T 8 j









I





El


I I
3ofl


p I


CM-'


I NOISE


SCCl0014 COMPENSATION


PATTERN


FIGURE I


I~00


LEVEL
















MICRONS


I I:


N-BUTYL


FIGURE 2


000


a~0O


-00

CM'


1 00


looD


PURE












































I 6I -Ioo


I 60


N-BUTYL


0.004 M


FIGURE 3


MICRONS


itOO


1 o


S400


/- I- r
"Y~r


* 1
















MICRONS


I I
~ItQG 3...


I #Oo


N-BUTYL CHLORO


FIGURE 4


1%,00
cm- I


Iava


PURE















































-~---I, I


Iloo
CM-t


N-BUTYL CHLORO


FIGURE 5


MICRONS


fVPkr\
V


~OOQ


raoc


0.006















MICRONS


I I


sow. ISSO


CM-


N-BUTYL DICHLORO

FIG URE


PURE


t1ac0













MICRONS


I I I


p4 /


I I -4--
S.c. ISc.


N-BUTYL DICHLORO


0. 006 M


FIGURE 7


CM- I
CM-I


.. ...... ........



















MICRONS
6 I TI
--- I --- -- I --- -- I-----i______


I -4---
5..~ IIbo


I I I


t"-oO


C M-




N-BUTYL CHLOROFLUORO


FIGURE 8


I o00


f co


PURE
















MICRONS


35
I I


i i
It --sq -


CM-'


N-BUTYL GHLOROFLUORO


0.005 M


FIGURE 9


I0












45

MICRONS


tooo 3000


N-BUTYL


SItoo IVM Itoo I40 0ee
CM-"


DIFLUORO


FIGURE 10


3S


PURE

















MICRONS


& ~ ~T4o''


I I


rAfi i oo l tc


CM-1


N-BUTYL DIFLUORO


0.009 M


FIGURE II


I o00


tooo Ao-x

















MICRONS


' I
.~oO


CM-I


N-BUTYL TRIFLUORO


FIGURE


I -4--
)8~o


tl.o 0


to0*


PURE











48
MICRONS


I3I Io Io I I I


CM-'


N-BUTYL


TRIFLUORO


0.004 M


FIGURE 13
















MICRONS
T


I I


I I~b'~ D100


N-DIBUTYL


FIGURE 14


CM-1


1000


PURE













50

MICRONS
6 7 9 q Tc


CM-,


N-DIBUTYL


0.005 M


FIGURE 15


3Oc~


S'8 I i i I I I
1300 16" 1*-o0 stolo cc

















MICRONS


S6 7


-oo B*L* soo* 1io.





N-DIBUTYL CHLORO


I I


PURE


FIGURE 16


I .oo


isoo.


GM-1




















MICRONS


I I


T 8
I I


9 g0
I I


Iow I 8o-
3000 o~ 180.


N-DIBUTYL CHLORO


FIGURE 17


/Io
, 6oo


I -'oo


I oo@


C M-'


0ooo0


0.007 M


I


I I












53

MICRONS


A,.6 3.
I ,I


I I


CM-I




N-DIBUTYL DICHLORO


FIGURE 18


I.


PURE














MICRONS


I I
*t~@oo ~.4Q


I I


CM-1


N-DIBUTYL DICHLORO


0.008 M


FIGURE 19

















MICRONS


ooo ,ooo ( goo itoo f *.i

CM-'






N-DIBUTYL CHLOROFLUORO


FIGURE 20


r .0


I Oo


PURE


' I I
















MICRONS


o o Ioo


CM~-


N-DIBUTYL CHLOkOFLUORO


0.007 M


FIGURE 21


I o a t* IA o Io


, 1


I I
















MICRONS


I I I I


i I I


C




N-DIBUTYL DIFLUORO


I .o








PURE


FIGURE 22


#OaO


o0e















MICRONS


I I


3..Q


CM-I


N-DIBUTYL DIFLUORO


0.007 M


FIGURE 23


I














'MICRONS


CM-,


N-DIBUTYL


TRIFLUORO


FIGURE 24


PURE


3.5































4000 .50"














60

MICRONS


a 6


I I II I I


i I I I
ooo .. laQ+ VID


l I' lO Ia
1~0o 't~ 'a.. lope


CM-'


N-DIBUTYL TRIFLUORO


0.006 M


FIGURE 25











CHAPTER IV

DISCUSSION

The most interesting general features of this research are the

clearly defined frequency shifts found upon systematic substitution

of halogen atoms upon the number two carbon atom of the acetamides

studied. Further, these shifts are not in random order, but may be

significantly related to the electronegativity of the substituent.

Discussions of the relative electronegativities of chlorine and

fluorine must take into account two very general phenomena. The

most obvious of these is the ability of the halogen atom to attract

electrons inductively, the I effect. On the basis of the electronic

screening of the nucleus, it follows that fluorine should possess the

stronger I effect. This is borne out by experiment. The second

effect is the resonance or mesomeric shift of electrons from the

valence shell of the halogen concerned into a coordinate bond. It

has been determined on the basis of a large amount of experimental

work that fluorine has a greater tendency to operate this / M shift

than has chlorine. Although Ingold has stated34 that the energies

Involved in the I and -AL M effects may be separated as the results

of two distinct phenomena, it is well to recall to mind that the effect

of any one halogen is unique. In other, words, pragmatically, it











62

must be recognized that there is only one effect, that of the replace-

ment of a hydrogen atom by a halogen atom.

In, the compounds under consideration, there is little or no

possibility for the operation of the / M effect. In order for the

halogen to shift a pair of electrons into the halogen-carbon bond,

there must be some mechanism for the removal of an electron pair

from the carbon atom. This follows unless pentavalent carbon is to

be assumed. If, on the other hand, two halogens are attached to the

same carbon atom, the / M shift of the one may be enchanced by

the I effect of the other. This actually would vitiate the I effect

with respect to the remainder of'the molecule.

Perhaps the most satisfactory quantitative data available for

a comparison of the effects of the halogens are the ionization con-

stants of substituted acids. Deductions concerning the inductive

effect introduced into the compounds investigated in the present study

are of the highest validity if related to the halogenated acetic acids.

The ionization constants of chloro-, dichloro-, and trichloro-acetic

acids are 1.396 x 10 5 5.5 x 10"2 and 0. 13, respectively. 29

Henne31 reports the ionization constants of the analogous fluoroacetic

acid series to be 2.17 x 10-3, 5.7 x 10-2 and 0.588 or 0.533. This

clearly demonstrates that the net inductive effect attributable to
fluorine is greater than that of chlorine.











63

In order to facilitate the presentation of the interpretation of

the data, a step out of logical order. will be made by first postu-

lating a configuration of the amides. This configuration will then

be shown to account for the infrared absorption spectra of the

N-monosubstituted and N-disubstituted amides in the region studied.

The interpretative work of Lenormant Is accepted with a

few reservations. It will be recalled that he proposed the N-mono-

substituted amides to be in a mesomeric state of the normally

written ketonic form and the corresponding excited state. He then

stated that the 6.40 micron (1562 cm-1) band in these compounds

was due to the stretch of the nitrogen-carbon partial double bond.

The fact that this band does not occur in the N-disubstituted aides

was not explained.

From the present works it is concluded that this band does

occur in the N-disubstituted amides, but that it is masked by acciden-

tal degeneracy with the carbonyl band. It is concluded that this band

and the carbonyl band are both found in the immediate neighborhood

of 6.00 microns (1667 cm-1).

It has been stated above that an effort has been made to

explain the absence from the spectra of the N-disubstituted amides

of the band in question on a steric basis. This was necessary since













it was recognized that the replacement of an amino hydrogen atom

by an electron releasing group would definitely increase the likeli-

hood of the assumption by the carbon-nitrogen link of some double

bond character. Furthermore, from the bond distances found in

acetamide, 62 some double bond character must be attributed even

to that carbon-nitrogen bond.

Rather strenuous objection may be raised to the idea of acci-

dental degeneracy in these compounds because of its generality.

It is obvious that there are certain minimal energy limits involved

in the assumption by heteroatomic systems of completely symme-

trized configurations. It is this limit which has prevented the

realization of such a state by the N-monosubstituted amides. The

substitution of an additional electron releasing group upon the nitro-

gen allows the system to cross this energy threshold and attain

the symmetrical condition described above.

The lowest energy state should be# according to the resonance

theory, that which is the most symmetrical with respect to the

electronegativities of the participating atoms. It is believed that

this state can be attained by the N-disubstituted amides and that

its realization is attested to by the absence of the 6.40 micron

(1562 cm"1) band.











65

Table VIII lists the data of Tables V and VI which are per-

tinent to the discussion of the N-monosubstituted amides. The B

band shift is that of the most prominent peak in the band. It should

be recalled that the N-butyl chloro compound was not run in the pure

state, but was mulled with carbon tetrachloride. For this reason,

the shifts shown for it are possibly not entirely comparable to

those of the other compounds. At the most, however, they are only

slightly lower than the true value.

TABLE VIII

A BAND AND B BAND DILUTION SHIFTS, N-BUTYL SERIES

Compound A Band Shift B Band Shift
(cm-1) (cm-1)

N-butyl 3SO 58

N-butyl chloro / 25 31

N-butyl dichloro 7 34 14

N-butyl chlorofluoro L 34 31

N-butyl difluoro / 34 29

N-butyl trifluoro / 30 41

The notable feature of Table VIII is the more or less central

position which the dichloro derivative assumes; that is, this com-

pound shows the greatest shift of the A band and the least shift of










66

the B band. There are no striking differences in the frequencies of

the A bands; however, Table V has shown that there is a steady drift

toward higher energies of both the pure and dilute A band absorp-

tions of the N-butyl amides. Assuming that this band is the carbonyl

stretch, the implication of these data is that the substituted halogens

exercise the expected I effect. This serves to inhibit the outward

shift of the electron cloud in the neighborhood of the oxygen atom.

The frequency of this absorption, in the case of the trifluoro deriv-

ative, is 1736 cm-1 (5.76 microns) in dilute solution. This is near

the high frequency limit, 1740 cm-1 (5. 75 microns), of the absorption

of the carbonyl band observed by Batuev 1 in dilute dioxane solutions

of the lower fatty acids. He found that the average shift of this band

is 70-84 cm"1 on dilution.

Table VI, on the other hand, shows that the net shift of the B

band is toward lower frequencies on dilution. It is seen that the

minimum shift is found in the case of the dichloro derivative; however,

the B band in the pure material occurs at lower energies than it

does in any other of the N-monosubstituted amides. The B bands of

the dilute, or unperturbed, compounds progressively increase in

frequency. Thus, there is a tendency on the part of the substituted

halogens to strengthen the bond with which the B band is associated,











67

but this does not hold true in the pure materials. Obviously, some

other effect must be sought in explanation of this observation.

The pure liquid N-monosubstituted amides are highly associ-

ated. This association involves the bonding of the amino hydrogen

to a donor atom such as the carbonyl oxygen, the amide nitrogen or

a fluorine. It is fairly evident, from the shift of the A band upon

dilution, that the oxygen atom has participated in a rather high

order of hydrogen bonding. A sensibly reverse phenomenon is

observed with the B band upon dilution. If this band is caused by

the absorption of a carbon-nitrogen bond of rather high double bond

character, it is obvious that a deformation of the electron shell

about the oxygen to a position away from the carbon-oxygen bond

would facilitate a shift of the electron density about the nitrogen

into the nitrogen-carbon bond. These tendencies are reinforced by

hydrogen bond formation. This is the situation which pertains in

the pure material. Upon dilution, the hydrogen bonding through

the oxygen is destroyed and is accompanied by a consequent shift

inward of the electron cloud about the oxygen. Thus, the observed

behavior of the B band upon dilution is that which would be expected

if the postulate above is correct.
The frequency of the B band in the pure materials is some -
what anomalous. Reference to Table VI shows that the most











68

prominent band tends to shift toward lower energies as the dichloro

compound is approached and then to higher energies upon further

substitution. Since the dilute solutions have been shown to behave

as expected, it is necessary to consider possible differences in

the nature of the association processes as substitution progresses.

Table V has shown that the oxygen atoms become less susceptible

to hydrogen bonding as the electronegativity of the adjacent group

increases. Examination of the frequencies of the unperturbed B

bands shows that this is also true with respect to the nitrogen, but

to a lesser degree. This fact calls attention to the possibility of

hydrogen bonding through the nitrogen atom. If this occurs, it

can be seen that the carbon-nitrogen bond is deprived of electrons,

thus becoming weaker.

The reasoning in the paragraph above is no longer applicable

upon encountering the chlorofluoro compound for there it is found

that the frequency of the B band in the pure compound is higher

than that in the pure dichloro amide, This trend continues as the

substituents become more electronegative. It will be noticed that

this tendency appears upon the introduction of fluorine into the series.

The extreme electronegativity of this element serves to increase the

double bond character of both the carbon-oxygen and the carbon-
nitrogen links. Simultaneously, the electronic density in the











69

neighborhood of the fluorine increases. These three effects pro-

mote hydrogen bonding through the fluorine rather than through

either the oxygen or the nitrogen. The occurrence of hydrogen

bonding of this type serves further to increase theprotonic nature

of the amino hydrogen, thus releasing more electrons for partici-

pation in the nitrogen-carbon band.

The scheme of association postulated above does not find

application in the discussion of the N-dibutyl amides. This is to

be expected in view of the absence of an amino hydrogen atom. In

Table V, it has been shown that the absorption frequency of the

carbonyl band shifts to higher values as the electronegativity of

the halogen substituent increases. These values are much lower

than those found for the unperturbed N-butyl derivatives. The be-

havior is that to be expected purely on the basis of the I effect.

On dilution, except for a general sharpening, there is no change

in this band for the first two members of the series.

An apparent anomaly is found upon examination of the dichloro

compound. The carbonyl band splits, one part shifting to higher

and the other part to lower frequencies. than that of the center of

the perturbed band. This effect is found to a smaller extent in the

chlorofluoro compound and even less in the difluoro. It is not











70

observed in the case of the N-dibutyl trifluoro amide. This behavior

must be explained in terms of the underlying concepts involved in

the hypothesis of accidental degeneracy in these compounds.

In the three compounds presenting this split, there exists a

rather highly activated hydrogen atom in the methyl group; There

is no question of the ability of this atom to participate in hydrogen

bonding through the oxygen or the nitrogen and, except in the first

case, through a fluorine Upon dilution, this intermolecular

effect is destroyed. This would strengthen the carbon-o:xygen bond

or the carbon-nittogen bond depending upon which atom was the

donor.

Sannie and Po remskii 61 as has been previously mentioned,

postulated the formation of an intramolecular hydrogen bond in the

disubstituted amides. This bond involves the carbonyl oxygen and

an alkyl hydrogen of one of the N-substituentsi This hydrogen bond

would be much weaker than one involving a hydrogen initially bound

to an atom rnore electronegative than carbon. An intramolecular

bridge would not be expected to be broken upon dilution. It symme-

trization initially exists in the presence of this intramolecular bond

and an intermolecular bond to the nitrogen, It will at least conceP-

tually, be destroyed upon dilution. The situation pertaining in











71

dilute solution is rather complex. The intramolecular bond still

exists, pulling electrons out of the carbon-oxygen bond and, through

the alkyl group, feeding the carbon-nitrogen bond. The electronega-

tive substituents are shifting the electronic cloud of the compound

toward themselves. If this occurs with symmetrization, there can

be no splitting of the 6.00 micron (1667 cm"^) band.

The net effect of dilution, then, is to strengthen the carbon-

nitrogen bond. Further, the perturbation on the nitrogen is removed.

This results in a sharpening of the band associated with the carbon-

nitrogen valence stretch. A symmetrical strengthening of the

oxygen-carbon and nitrogen-carbon bonds must involve a weakening

of the intramolecular hydrogen bond. On the other hand, a strength-

ening of the carbon-nitrogen bond alone tends to increase the

stability of the internal hydrogen bond. In view of the phenomenon

observed, it must be concluded that the least energy configuration

for these compounds demands the existence of this hydrogen bond;

thus, a non-symmetrized state. From the intensities of the portions

of the split band, it can be concluded that the higher energy contri-

bution is that of the carbon-nitrogen bond.

It must be inferred from the argument above that the internal

hydrogen bond exists in all six of the dibutyl derivatives. The











72

absence of the split band, even in dilute solution, from the spectra

of some of the dibutyl compounds tends to discount the preceding

argument. In the cases of the non-halogenated and rmonochloro

derivatives, this can be explained by the fact that insufficient elec-

tronegative stress is placed in the molecule to effect the separation.

That is, the hydrogen bond is stabilized internally at very nearly

the proper energy to effect complete overlap of the carbon-nitrogen

and carbon-oxygen Stretches. Consideration of the trifluoro deri-

vative shows that the extremely high electronegativity of the tri-

ftuoromethyl group should be more than sufficient to split this

band. Observation of the spectra shows that the split is observed

to its greatest extent in the case of the dichloro compound and that

it becomes smaller as the electronegativities of the substituents

increase. This implies that at lower electronegativities there is

a mean displacement of the electron cloud which stabilizes the

hydrogen bond. Further increase in the strength of the electroneg-

ative center forces the hydrogen bond to break in favor of the

greater electron availability introduced in the unbound oxygen.

This once again leads to the almost complete symmetrization of

the nitrogen-carbon and oxygen-carbon links in the case of the

trifluoro derivative.











73

Inspection of the spectra of the pure N-disubstituted amides

shows that these bands are amply wide for the concealment of two

peaks split on the order indicated by dilution. This natural breadth

is intensified by intermolecular interactions of a lower order of

energy than that of hydrogen bonding.

Reference has been made to hydrogen bonding possibilities

through the fluorine atoms in the compounds studied. Table VII

shows that there are definite shifts in the carbon-fluorine valence

stretch frequencies upon dilution. These shifts are of greater mag-

nitude in the N-butyl compounds. This implies that, even though

they are activated, the alpha hydrogen atoms are not as electroposi-

tive as is an amino hydrogen. Sutherland states that if the shift

in frequency is approximately three percent of the frequency of the

band in question, it may be assumed that hydrogen bonding has

occurred. Although the shifts being referred to here are all less

than this value, it is believed that they stand in evidence of hydrogen

bonding. Sutherland, in his discussion, placed emphasis upon

shifts of absorption frequencies of bonds of the acceptor atom with-

out reference to the donor atom.

The general weakness of the shifts in the chlorofluoro com-

pounds and the absence of a shift in the N-dibutyl trifluoro amide










74

are of particular significance. The latter, having no acidic hydro-

gens, would not be expected to participate in intermolecular bonds.

In the former pair, the degree of activation of both the halogens and

the methyl hydrogen are of rather low order.

Several references to hydrogen bonding through fluorine are

available. Though most of these involve the hydrogen fluoride case,

several are concerned with other molecules. Copley, Zellhoefer

and Marvel15 investigated bonding of fluorine-activated hydrogens

to amides. They found that these hydrogens are bonded to the nitro-

gen or oxygen of the disubstituted amides only. The conclusion

reached to that no matter how highly activated it may be. a hydrogen

bound to a carbon atom is not capable of forming a hydrogen bond in

preference to the bonds usually occurring in either the amides or in

their N-monosubstituted derivatives. This study was based upon

the solubility of dichlorofluoro methane in several amides. Marvel,

Copley and Ginsberg52 found that the heat of mixing of benzotrifluo-

ride in N-dimethyl acetamide is much greater than that of benzo-

trichloride in the same solvent. On the other hand, it was found

that the heats of mixing of the two solutes in acetone are the same.

This difference in energy was interpreted as being due to the fact

that hydrogen bonding initially exists in benzotrifluoride, the











75

p-hydrogen acting as the acceptor atom, From cryEtallographic

data, 64 it has long been know that a hydrogen bridge Involving flu-

orine as the donor atom exists in the ammonium fluoride crystal.

Table VII shows that the stretching modes of the carbon-

fluorine link have been found in the 8.28-9.68 micron (1208-1033

cm-") range. It is noted that characteristic bands of this bond are

found at higher frequencies as the number of fluorine atoms present

increases. This phenomenon is analogous to the Scanlan-W'Vrhurst

Effect63 which is described as the shortening of the carbon-halogen

bonds as the number of halogen atoms upon a given carbon atom

increases. This holds true in molecules in which resonance occurs

between covalent and ionic forms. It may be added that there is a

high probability of finding this type of resonance in any fluorine-

carbon bond because of the large difference in electronegativity

between the two elements..

The observed shifts of the carbon-fluorine frequencies are all

in a negative direction upon the breaking of the hydrogen bridges.

This must be explained by stating that the shift of electrons out of

the bond through the agency of hydrogen bonding is more than bal-

anced by the increased availability of electrons from the carbon-

nitrogen bond.











76

An interesting generality mentioned by Buswell, RodeDuuL ..M.

Roy14 with reference to hydrogen bonding has been stated by

Venkateswaran. The Rule of Venkateswaran relates the acidity

of a hydrogen participating in a hydrogen bond to the frequency shift

observed upon the formation of that bond; that is, the more acid the

hydrogen, the greater the shift toward lower frequencies and the

broader the resulting absorption band. Reference to the spectra

reproduced in this dissertation shows that the bands of the pure com-

pounds listed in Table III definitely differ in broadness. In each

case, the band occurring in the 3.05 micron (3279 cm-1) region is

much broader then the neighboring band at 3.25 microns (3077 cm'l).

If the Rule of Ven-kateswaren is applicable, its implication is obvious.

It is interesting to observe that there is some progression in

the frequencies of the two perturbed bands of the nitrogen-hydrogen

stretch as was seen in the case of the perturbed B bands. Although

the differences in wavelengths are on the order of magnitude of the

error to be expected, it is assumed that they are significant because

of their regular trend. From the lower frequency of the perturbed

nitrogen-hydrogen valence stretch in the case of the dichloro com-

pound, one can deduce that the amino hydrogen is more protonic

in this instance than it is in the other compounds. Thus, the approach











77

of a hydrogen atom to the nitrogen results in an average increase

in the protonic character of the amino hydrogen atoms.

An important fact to be borne in mind in studying the bands

in this region is that, energy-wise, the wavelength plot utilized

presents a false band width when compared to the bands in the

lower frequency regions. In other words, the actual energy spread

of the bands corresponding to the nitrogen-hydrogen stretch is far

larger than that of the A and B bands. For this reason, the large

number of possible bonding modes which exist are masked by the

method of recording inherent in the instrument used.

There is one salient fact in the data of Table Ill. This is

that the shifts of the nitrogen-hydrogen stretch bands in the fluoro

compounds tend to be smaller than those of the other compounds.

The implication is that hydrogen bonds through fluorine are weaker

than those through oxygen cr nitrogen in the present cases. Exam-

ination of the cyclic dimers which would be formed through fluorine

as contrasted to those through the other two donor atoms yields a

clue to the reason for this. A dimer through amino hydrogen bridges

to fluorine involves a ten-membered ring, but the dimer between the

same acceptor and the carbonyl oxygen is only an eight-membered

ring. Aside from dubious remarks which might be made regarding











78

steric effects, it can be seen that a quasi-resonance stabilization

is more probable the smaller the ring. This comment, of course,

does not refer to conjugated double bond systems.

Several conflicting opinions are available regarding the possi-

bility of distinguishing between amino hydrogen bonds to nitrogen

and those to oxygen with the resolving power of the instrument

being used in this research. Stanford and Gordy reported ex -

tremely small shifts of the absorption frequency of the acetylenic

hydrogen in phenyl acetylene when it was dissolved in either di-

methyl acetamide, dimethyl formamide or acetone. Corresponding

to heat of mixing calculations, they found that a smaller perturba-

tion was introduced in acetone solution than was present in the

solutions of the two amides, From this, they concluded that the

hydrogen bond to nitrogen is stronger than that to oxygen. Of

course, these data might just as easily have been interpreted as a

measurement of the relative electronegativities of the various

carbonyl groups involved. Stanford and Gordy assumed that the

amide nitrogen is more basic than is the amide oxygen. That this

is not necessarily the case has been demonstrated in this research.
4
Anzilotti and Curran determined the percentage of o-fluorophenol

molecules possessing internal hydrogen bonds from spectral and











79

dielectric constant measurements. Although the calculated elec-

tron density upon both atoms is the same, Curran17 found that the

oxygens in catechol were more strongly hydrogen-bound than was

the fluorine in the compound mentioned above. Buswell, Rodebush

and Roy14 had previously concluded that the hydrogen bonds to

nitrogen and oxygen in the amides are indistinguishable.

Upon dilution, the nitrogen-hydrogen stretching modes become

almost completely unperturbed. This fact follows from the disap-

pearance of the two important bands discussed above and the

appearance of a band at 2. 92 microns (3425 cm-1). There is an

apparent tendency for the lower frequency band to disappear first

upon dilution. Initially, this band is less intense than is the higher

frequency band. This is due to the fact that, though the formation

of a ring tends to stabilize the system as the magnitudes of the

frequency shifts show, the probability of obtaining the orientation

necessary to the formation of a ring is only one-half that of form-

ing a linear dimer, An explanation which does not consider the

entropy change upon dimerization is that resonance stabilization in

a ring system reduces the effective dipole moments of the absorbing

bonds, thus decreasing the intensity of the absorption.











80

It is interesting to note that the unperturbed nitrogen-hydrogen

stretch band is present in the pure liquid N-butyl difluoro amide.

As has been mentioned, this compound is expected to have the most

acidic hydrogen of the series other than the amino hydrogen. The

presence of the 2.92 micron (3425 cm-1) band is an indication of the

fact that a fraction of the available donor atoms is blocked with

respect to the amino nitrogen because of bonding through the active

methyl hydrogen.

The question of keto-enol tautomerism has not been satisfac-

torily settled. Because of the broadness of the bands in the three

micron region, one can postulate that perturbed hydroxyl stretches

actually exist. On the other hand, in extremely dilute solutions,

there is no evidence whatsoever of the existence of a true hydroxyl

stretching mode. This limits the existence of the enolic form to

the cases in which there is actual association. It may be stated that

upon the formation of a hydrogen bridge through an oxygen, there

is some doubt as to the position of the proton; that is, that the proton

resonates between the ironically and covalently bound states. It is

certain that the classical tautomeric mixture does not exist in dilute

solution.

It has been mentioned before that the cyclopeptides do not
absorb radiation in the range of the B band.51 Diketopiperazine,











81

during basic hydrolysis, gradually begins to absorb in the 6.40

micron (1562 cm"1) region. The band at 5.95 microns (1681 cm"1)

disappears and another band appears at 6.08 microns (1645 cm-1).

It was mentioned in Chapter I that the carbon-nitrogen bond length

in diketopiperazine is shorter than that of the corresponding bond in

acetamide. This implies that the carbon-nitrogen bond in the former

compound is of greater double bond character than it is in the latter.

Several resonance structures may be drawn for this compound. The

ordinary structure may be coupled with an excited state in which

both nitrogens are doubly bound in the ring. There is an alternative

hybrid which may lie between the two equivalent structures in which

only one of the amide -like groups is excited at a time. This is

roughly analogous to the Kekule structures written for benzene.

Resonance between two such structures would be expected to be com-

plete, involving a high stabilization energy.

It must be concluded, after consideration of the pronounced

possibilities for resonance hybridization, that the diketopiperazine

molecule behaves in a fashion analogous to that of the N, N-disubsti-

tuted acetamides. Thus, the band observed at 5.95 microns (1681

cm^-) is characteristic of both the carbon-nitrogen and carbon-

oxygen bonds. Upon attack by sodium hydroxide, this resonance is











82

destroyed and the ring is split. The two bands then appearing may

be identified as corresponding to the A and B bands in the N-nmoro-

substituted amides.

The linear peptides behave in a fashion identical to that of the

corresponding amides. Lenormant51 states that a linear polyamide

of the nylon type absorbs very strongly at 6.36 microns (1572 cm"1)

and 6.0 microns (1667 cm-1). Upon methylation of the nitrogen,

only the very strong absorption at 6.05 microns (1653 em"I) appears.

The same author, with reference to the lactams, states that

they do not absorb in the B band range. The lactam of 6-amino-

hexanoic acid absorbs at six microns. Dilution data are not avail-

able for this compound. The occurrence of a symmetrized system

is hardly to be expected in this compound. A possible explanation

for the absence of the B band is that the carbon-nitrogen bond is

virtually one-hundred percent single bond in character. A reason

for this is not advanced. Room temperature treatment with one-tenth

normal sodium hydroxide followed by frequent checks of the spectra

would hardly solve this problem. The hydrolysis product is the

salt of an amino acid. The carbonyl group is too far removed to

enter into conjugation with the 6-amino group.











83

Succinimide poses a problem similar to that of caprolactam.1L

The B band is not found in the spectrum of this compound. This can

be explained by observation of the improbability of the fomation of

a double bond from the nitrogen in the presence of the competing

groups.

Lenormant51 has reproduced the spectra of several N-monosub-

stituted acetamide compounds. These include the sodium, mercury.

bromine and ethyl derivatives. Sodium acetamide does not posses

a band absorbing in the A region, but does absorb 6.37 microns

(1570 cnim). The di-acetamide salt of mercury absorbs at 6.23

and 6.33 microns (1605 and 1580 cm'1). N-bromo acetamide does

not have a band in the B region, but absorbs strongly at 6.06 microns

(1650 cm'l), As would be expected, N-ethyl acetamide possess

both A and B bands, absorbing at 6.00 microns (1667 cm-1) and

6.39 microns (1565 cm-l1)#

It has already been stated that the reason for the presence of

the A and B bands in the spectra of the N-monosubstituted amides is

that there is insufficient electronic density in the neighborhood of

the nitrogen atom to allow the symmetrization of the carbon-nitrogen

and carbon-oxygen bonds. The absence of the A band from the

spectrum of the sodium salt of acetamide implies that the carbonyl











84

bond does not exist in this compound. Unfortunately, these salts

are not totally unambiguous in structure. The reaction of metallic

sodium with acetamide is accompanied by a liberation of hydrogen.

One would deduce from this fact alone that a negative charge

resides upon the nitrogen. Further, the reaction of the sodium

salt with an alkyl iodide yields the N-alkyl acetamide, A most

elementary consideration of the electronegativities of nitrogen and

oxygen, on the other hand, leads directly to the proposal that the

structure of the salt involves a double bond from the nitrogen to

the carbon with the negative charge of the ion virtually localized

upon the oxygen. This latter fact and the spectral data seem to be

in agreement. In view of the tendency of the compounds of nitrogen

to undergo rearrangements in their reactions, one must suppose

that such a mechanism pertains in the formation of the N-alkyl

amides by the method described.

It is established, then, to a reasonable degree of certainty,

that the B band is characteristic of the valence stretch of nitrogen

essentially doubly bound to carbon. LenormantSI has also repro-

duced the spectrum of ethyl imido acetate. It shows only one absorp-

tion in the region under discussion. This band occurs at 5.95

microns (1681 cm-1). This must be attributed to the










85

carbon-nitrogen double bond. It will be recalled that a similar

deduction was made concerning bands occurring in the neighborhood

of 5.93 microns (1686 cm"1) in the spectra of certain of the dilute

N-disubstituted amides studied in the present investigation. The

reason for the higher frequency of absorption of this bond in the Imido

acetate as contrasted to that of the amido-anion of the sodium salt

should be obvious. The electronic density in the neighborhood of the

oxygen is reduced considerably by the reaction of the latter with an

ethyl carbonium ion. This inductively strengthens the nitrogen-car-

bon double bond with the resulting higher absorption frequency.

There seems to be little doubt; as to the structure of the bromo

derivative of acetamide. The studies of the Hofmann degradation of

amides seem to indicate that the N-bromo configuration is the proper

one. This would appear to be the more logical in view of the electro-

negativities of the elements involved. In acetamide itself, Davies

and Hallam20 found that the carbonyl band occurs at 1695 cm-1 (5.90

microns) but that upon dilution it shifts to 1714 cm-1 (5.80 microns).

Since Lenormant did not specify the conditions under which the spec -

tra referred to above were taken, one must presume that they are of

the pure compounds. The absorption of N-bromo acetamide .s far

displaced from that of acetamide. This shift to lower energies











86

appears to be anomalous in the light of the I effect attributed to

bromine; however, the effect of association in the pure compound

has not been considered. In this instance, the amino hydrogen is

extremely acidic. It can form a stable hydrogen bridge with the

ox-ygen of another amide. Further stabilization will occur with

cyclization. This bridge will tend to weaken the absorption fre-

quency of the carbonyl bond. It can be surmised that this weakening

effect predominates over the I effect of the bromine because the

former is not spatially insulated. Further, the inductive transmis-

sion of charge to the bromine atom is probably more effective over

the oxygen-hydrogen-nitrogen path than it is over the oxygen-carbon-

nitrogen path. The absence of the B band is self-explanatory.

The absorption spectrum of the di-acetamide salt of mercury

is produced by a structure which is quite as mysterious as is that

of the sodium salt of acetamide. This former salt is prepared by

fusion of acetamide with mercuric oxide and is accompanied by a

splitting out of water. This salt can occur in one of two possible

configurations or a combination of these. The first involves ionic

bonding through the oxygen and the second, covalent bonding through

the nitrogen. This deduction is based on the nature of the nitrites of

silver and sodium.33 The former, in the solid state, appears to











87

exist with the metal bound to the nitrogen, and the latter, with the

metal adjacent to the oxygen. It is suspected that the silver salt

is actually covalent in nature. Consideration of the behavior of mer-

cury leads to the plausibility of the suggested modes of bonding in

the di-acetamide salt of that element. The existence of rather dis -

placed A and B bands implies that a covalent salt of mercury exists.

In this condition, symmetrization is approached, but is not attained

because of competition of the two amide groups for the electrons of

mercury.

In contract to succinimide which was discussed previously,

Lenormant51 found that IT, N-dibenzoyl hydrazine absorbs radiation

at 6.06 and 6.48 microns (1650 and 1543 cm"1). In this compound,

there is a tendency for the phenyl rings to act as an electron source

in the formation of negatively charged oxygen atoms and as a sink

with respect to this excited state. Thus, the carbonyl bond would

be expected to absorb at rather low energies. Further, there is the

possibility that as it acts as a sink, the phenyl ring can induce the

nitrogen to shift a pair of electrons partially into its bond with the

adjacent carbon. This accounts for the presence of the B band and

its rather low frequency.











88

Following the suggestion of Richards and Thompson, 59

Lenormant51 N-deuterated certain alkyl acetamides. As was men-

tioned in Chapter I1 the B band was shifted, but not in the proper

ratio. It is true that the reduced mass of the deuterated system

differs from that of the hydrogenated system. From the simplest

two-body approximation in which the absorption frequency is inverse-

ly proportional to the square root of the reduced mass, it is evident

that the system behaves properly. This, of course. assumes that

there is no change in the force constant, as there rightfully should

not be. Because of the quantitative inapplicability of the approximate

equation, no calculation of the shift is made on this basis.

It will be interesting to determine if the slightly divergent

series of compounds prepared by Cromwell, Miller, Johnson, Frank

and Wallace16 behave as would be predicted by the present postu-

lates. It should be recalled from Chapter I that the B band is

found in certain of the alpha-amino, alpha-beta unsaturated ketones.

This occurs in spite of the fact that the amino groups are disubstituted.

In several cases, an absorption characteristic of the double bond is

not observed. This tends to throw suspicion upon the analogy which

has been drawn with respect to the B band. The probable explanation

is that, the carbonyl bond and the carbon-carbon double bond being











89

conjugated, both are shifted to lower frequencies than would be

expected. This series of compounds, then, stands largely as an

externally consistent explanation of the lowered absorption frequency

of both the carbonyl bond and the nitrogen-carbon double bond. Here,

symmetrization is not possible because of the intervention of a

carbon-carbon single bond.

Much discussion of the relationship between the shifts of ab-

sorption frequencies of bonds through association and the chemical

heats of formation of hydrogen bonds has appeared in the literature

Several excellent reviews exist; however, the most recent and com-

plete is that by Davies. 19 It is the consensus of opinion that no

simple linear relationship exists between these shifts and the heats

of formation of hydrogen bonds. Davies has mentioned that numerous

forces must operate in the formation of a hydrogen bond. Among

these are dispersion forces and dipole-dipole forces. It is seen that

neither of these effects would necessarily be reflected in absorption

frequency shifts.

The largest frequency shifts observed in the present research

were on the order of 360 cm-1. This is equivalent to I. 03 kcal./bond-

mole. Davies and Hallam, 20 who proposed a cyclic trimer for the

associated forms of acetamide liquid, state that such a structure











90

would involve a heat of formation of about 1.1 kcal./bond-mole.

They report, however, that a non-spectroscopic determination of

the energy involved in hydrogen bond formation in aniline-benzo-

phenone mixtures yielded a value of 2.0 kcal./bond-mole. The heat

of formation of the formanilide dimer is reported to be Z. 9 kcal./

bond-mole.

A more nearly comparable case, that of the acetic acid dimer,

is reported by Pauling55 to involve an energy of 8. Z kcal./bond-

mole. Further discussion by Pauling indicates that the energy of

the hydrogen bond formed in ammonium fluoride crystals is about

5 kcal./bond-mole. It is seen that poor agreement is reached in the

spectroscopic case.

There seems to have been no discussion of the part which the

shift of the absorption frequency of the bond involving the donor atom

must play in the calculation of bond energies. If one were to con-

clude that these shifts are to be added, calculation of the energies

of hydrogen bonding from the present data would range from 0.54

kcal./bond-mole in the linear diners to 1.1 kcal./bond-mole in the

cyclic dimers. One might accept these values except for the very

unfavorable comparison to that mentioned above in the case of

acetic acid. It would seem that cyclization would, in the present











91

case, introduce sufficient stabilization to double or treble the bond

energies observed in the aniline-benzophenone dimer. V'irtz has

calculated that the resonance stabilization of hydrogen bonds of

amides is about 1 kcal. /bond-mole.74

Davies and Sutherland21 have reported the difference in the

frequencies of the hydroxyl stretch and the carbonyl stretch between

monomeric and dimeric acetic acid to be 448 cm-1 and 67 em-,1

respectively. The average shifts observed in the present research

are on the order of 360 cm-1 and 35 cm 1 for analogous bands.

Using as a basis the heat of association of acetic acid reported above

and the ratio of the major shift reported for acetic acid to the aver-

age major shift of the N-monosubstituted acetamides, it is calcu-

lated that the average heat of association of the latter compounds is

6.5 kcal. /bond-mole.

Badger and Bauer and, later, Badger, utilized a similar

method to obtain spectrographically the heats of association of

several alchols. It was reported that a simple graphical interpola-

tion method yielded values with an uncertainty of 500 cal. /bond-

mole. Although there is partial justification of the use of this

method in the present case, it is not to be taken as an absolutely
reliable measurement. Unfortunately, too little is known of the

more subtle forces operating between molecules.











92

It is believed that the hypothesis of .ccidental degeneracy

stated at the bcginnin-' of this chapter has been maintained in the

presence of experimental data obtained both in the present research

and in investigations reported in the literature. Though the expla-

nations at times seem to be rather devious, it must be observed

that the phenomena dealt with here are by no means simple and

have for many years resisted reduction to mathematical terms.











CHAPTER V

SUMMARY

1. It has been demonstrated on the basis of the present

research and in accord with previous conclusions that the 6.40

micron (1562 cm-1) band of the N-monosubstituted amides is attrib-

utable to the stretch of the carbon-nitrogen partial double bond.

2. Previously unreported absorptions have been found in the

spectra of the N, N-disubstituted acetamides studied. These bands

occur at 2.91 microns (3436 cm-1). 3.04 microns (3289 cm-"1) and

5.70 microns (1754 cm-1). The former bands are considered to be

overtone bands, but the latter is unclassified.

3. It is postulated that the absence of the 6.40 micron

(1562 em-1) band from the spectra of the N,N-disubstituted amides

is explained by the accidental degeneracy of this band and the car-

bonyl stretch band. This is the result of a nearly complete symme-

trization of the bond energies of the amide group.

4. By means of the shifts of absorption frequency of the

carbonyl bond upon halogenation of the N-substituted amides studied,

the fluorine atom has been shown to exert a greater electronegative

inductive effect when bound to carbon than does chlorine. This











94

is in agreement with previous determinations of the magnitude and

direction of the inductive effect.

5. It has been shown that hydrogen bonding exists in the

N, N-disubstituted aides in the case in which a halogen-activated

hydrogen atom is present. Furthermore, it is found that hydrogen

bonding in the compounds studied occurs through oxygen, nitrogen

and fluorine atorms.

6. A splitting of the 6.00 micron (1667 cm-1) band of N-dibutyl

dichloro, N-dibutyl chlorofluoro and N-dibutyl difluoro compounds

has been observed upon dilution.











BIBLIOGRAPHY


1. Aelion and Lenormant, Compt. rend. 224, 904 (1947).

2. American Petroleum Institute Research Project 44, "Catalog of

Selected Infrared Absorption Spectrograms", National Bureau of

Standards, Washington, D. C.

3. Anauthakrishnan, Proc. Indian Acad. Sci. SA, 200 (1937).

4. Anzllotti and Curran, J. Am. Chem. Soc. 65, 607 (1943).

5. Badger, J. Chem. Phys. 8, 288 (1940).

6. and Bauer, ibid., 5, 839 (1937).

7. Barnes and Bonner, J. Chem. Education 14, 564 (1937).

8. ibid., 15. 25 (1938).

9. R. B., Gore, R. C., Liddel, U. and Williams, V. Z..

"Infrared Spectroscopy"', Reinhold Publishing Co., New York,

1944.

10. Bates and Hobbs, J. Arn. Chem. Soc. 73, 2151 (1951).

11. Batuev, Compt. rend. acad. sci. U. R. S. S. 53, 507 (1946).

12. Buswell, Downing and Rodebush, J. Am. Chem. Soc. 62, 2759

(1940).

13. Krebs and Rodebush, J. Phys. Chem. 44, 1126 (1940).

14. Rodebush and Roy, J. Am. Chem. Soc. 60, 2444 (1938).

15. Copely, Zellhoefer and marvel, ibid., 60, 266 (1938).




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