Pa-dmp
4 California Association of Chemistry Teachers
George C. Pimentell University of California Berkeley
Infrared Spectroscopy: a Chemist's Tool
Twenty-five years ago few chemists could interpret an infrared spectrum and even rarer were those equipped to record one. Even for the spectroscopist this spectral region presented difficult experimental problems. Today, however, the importance of this type of spectroscopy is such that every chemist should know its capabilities. Indeed, a t the Uuiversity of California, every undergraduate chemistry major has the opportunity to record and interpret infrared spect,ra in two laboratory courses. I t is the purpose of this paper to explain why this branch of spectroscopy has assumed such a role of prestige in chemistry. Applications to chemistry will he enumerated and examples will be cited which have appeared in recent literature. A number of these problems have been selected from research conducted in the Berkeley laboratories, but the reader should nct let this provincialism obscure the wide utilization of this technique. What Is Infrared Spectroscopy?
Spectroscopy is, of course, the study of the interaction of light with matter. I n absorption spectroscopy, light of a known frequency and intensity is passed through a sample and the intensity attenuation is measured. A record of the attenuation (absorption) as a function of frequency or wavelength is called a spectrum. The term "infrared" means "beyond the red" and identifies th? spectral region of wavelengths exceeding about 7500 A, a region in which the human eye is not sensitive. This region is called, colloquially, the "heat" end of the spectrum because these frequencies, when absorbed, raise the temperature of the absorber. Two convenient devices for demonstrating this behavior are a thermocouple and a fingertip. The wavelengths in the infrared are roughly in the range 1O4-10B A. For convenience, the units are almost always expressed in microns (1 p = l O V mm), so that the spectral range of interest to us is 1 to 100 p. Even more convenient than wavelength, though, is the frequency, o, always expressed in units of "reciprocal centimeters" or "cm-I." The frequency in em-' is obtained by dividing u, the frequency in cycles per second, by c, the velocity of light.2
" = -.c '
cm-1 = cycles per second cm per second
'Based upon a talk presented at the 1959 Summer Conference of the Pacific Southwest Association of Chemistry Teachers. Sugust 17, 1959, Asilamar, California
The spectral range in these units is 10,000 to 100 em-'. The energy equivalent of a quantum of light is fixed by its frequency through the Planck relationship E = hu. By this proportionality, we see that the low frequencies of the infrared region can cause only small excitations. The energy absorbed is not sufficient to break chemical bonds or alter molecular configurations. The type of molecular excitation which can occur is the initiation of vibrational movements of the atoms to and fro about their equilibrium positions. A useful and accurate picture of these movements is provided by a molecular model assembled from wooden halls connected by springs in the appropriate spatial arrangement. The resonant vibrations of such a "ball and spring" model are exactly analogous to the movements of a molecule which has absorbed a quantum of infrared light. The frequencies at which resonances occur in the model are analogous to the characteristic frequencies a h r l ed by a molecule, as revealed by its spectrum. Figure 1 expresses these ideas schematically. The hall arid spring model represents carbon dioxide and the sequence of pictures suggests the interaction of the molecule with light. In the first sequence, a quantum of light of frequency 2349 em-' is absorbed, exciting a characteristic resonant vibrational motion of CO?. In this motion the carbon atom moves along the oxygenoxygen line in opposition to the movements of the oxygen atoms. This motion is called the "asymmetric stretching moden-the word "stretchmg" indicating that the bond lengths are periodically changing and the word "asymmetric" indicating that one bond hecomes longer when the other becomes shorter. Returning to Figure 1, we see that the second quantum is not absorbed because there is no characteristic vibrational motion of COz with this freauencv. The third lSemanticists argue that 5 should not be called a "frequency" and they prefer an alternate term, "wave number.'' In fact, the units cm-I are oommonly called "wave numbers." These 6emantioists are indeed correct, since the units of u are not those of a frequency, oycles per second. Y e t i conveys the information contained in the frequency (via.the constant conversion factor, e ) and spectroscopists regularly refer to 5 as "the frequency." It would be misleading to suggest that the term "freauencv" . . is used only in reference to F y c ~ e s G rsecond. Quite the opposite is true i n infrared, visible, and ultraviolet speotroseopy, spectral regions in which the units are generally cm-'and the units cycles per second are rarely used. It might be added that there has been a recent move toward the adoption of the name "Kayser" for "om-'." This unit has the disadvantage of replacing a self-explanatory unit, em-', by a man's nrtmc. The term "Kavser" is now virtuallv out of use in United States journals of spee&oseopy. Volume 37, Number 12, December 1960
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quantum is again absorbed but this time exciting a different resonant vibrational motion. This motion is called a "bending mode" for obvious reasons and its frequency is 667 em-'. The two vibrations of C 0 2 pictured in Figure 1 are called fundamental modes. A non-linear molecule made up of N atoms has 3N-6 such fundamental modes. (A linear molecule has 3N-5). Thus we find that small molecules, like COz and CH4, have relatively simple spectra whereas complicated molecules, like acetylsalicylic acid (aspirin) and ascorbic acid (vitamin C), have complex infrared spectra. Nevertheless we shall see that the spectra of all of these molecules are of value to the chemist. In anticipation of the more detailed discussion to follow, a consideration of the "ball and spring" model reveals why infrared spectroscopy is such a fruitful tool in the hands of a chemist. What iixes the resonant frequencies of the "ball and spring" model? The masses of the balls, the strengths of the springs, and their geometrical arrangement, of course. Extrapolating to the molecular scale, then, it is the masses o i the atoms, the strengths of the chemical bonds, and the molecular architecture which determine the infrared frequencies absorbed by a molecule. The chemist knows the masses of the atoms and he cares a great deal about learning the strengths of chemical bonds and the spatial arrangement of the atoms. Clearly the vibrational spectrum warrants further consideration.
deed this is true--almost anything with chemical bonds absorbs infrared light. In fact the only molecules that do not absorb are those like O2and N2which, by reason of molecular symmetry, do not possess a dipole moment in any degree of vibrational excitation. To put it crudely, there is no dipole moment rhange associated with vibrational excitation, hence no electrical "handle" for the light to grab hold of. But all unsymmetrical diatomic and all larger molecules do have such a "handle" and do ahsorb infrared light. As a corollary, we see that it will he difficult to find substances which do not absorb anywhere in the infrared region. The only materials which are transparent in quite thick layers are ionic salts. These solids do not possess chemical bonds of the usual type and consequently their vibrational absorptions are in a different spectral region. This is fortunate for the spectroscopist who needs transparent window and prism materials (glass is opaque in the infrared!). Conventional infrared spectroscopy is based on the use of prisms made of NaCI, KBr, CaF2, LiF, CsBr, and CsI. H o w Are Infrared Spectra Recorded?
To record an infrared spectrum, the chemist must have a spectrometer. Basically any spectrometer consists of a light source, a dispersing element (to separate frequencies), a detector, and focusing optics. These elements are shown in a schematic way in Figure 2. In this simple spectrometer, light from the source, S, is focused by mirror MI onto the jaws of a narrow slit, S:. This slit may be as narrow as 0.01 mm or as wide as 1 mm, depending upon the optical system and its resolving power. The light which passes through S, is focused into parallel light by mirror M? prior to passage through the prism, P. The prism is usually a large, beautifully polished, single crystal of sodium
M3 Figure 2.
Figure 1. A schematic view d the excitation of vibrotionr of corbon dioxide by infrared light.
What Substances Absorb Infrared Light?
The discussion just given suggests that many substances should absorb somewhere in the infrared. In652
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SZ
D
M4
A schematic optical path of a n infrared qxctrorneter.
chloride. This material has excellent optical properties (dispersion and transparency) in the spectral region which interests the chemist. Now as the infrared light leaves the prism, each frequency is refracted through a unique angle so that mirror Ms focuses only a narrow range of frequencies onto the second slit, S2. This selected frequency (now almost monochromatic light) is focused by M p onto the detector, D. To scan a spectral region, Ma can be rotated slowly, so as to alter continuously the frequency which falls on the slit, S2. The detector response as a function of angle of rotation of Ms furnishes a spectrum. Calibration permits a
conversion of the angle of rotation to a frequency scale. If, now, an absorbing sample is inserted somewhere in the optical path, say a t X, the detector response will be modified in accordance with the absorption of the sample. A comparison of detector response with and without the sample as a function of the Ma mirror angle gives the desired information, that is, the infrared frequencies absorbed by the sample. Though it pays to understand the operation of this important laboratory tool, it is fortunate that each research ~ o r k e rdoes not have to assemble his own instrument. The revolutionary increase in the use of t,his method of study has been possible because of the availability at moderate cost of excellent commercial spectrometers. Easily operated and extremely useful instruments can be purchased for a few thousand dollars and the most versatile research instruments are priced in the range $10,000 to $20,000. Most of these commercial instruments use salt prisms, though for the highest resolving power gratings are employed. The light source can be a piece of carhorundnm, ceramic, or tungsten ribbon heated by an electric current to a temperature in the range 10002000°C. The detector is usually a thermocouple though there are photoconductive and pneumatic devices which give greater detection sensitivities in return for a few experimental headaches. The focusing optics are always mirrors (front surface) since glass lenses are opaque and salt lenses are difficult to grind. All of the commercial models include recording devices which provide an inked record of the light transmitted by any sample placed in the optical path over the spectral region which can be scanned by the instrument. The usual spectral scan requires about 15 to 30 min. Solids, liquids, and gases are all readily studied-a 10-cm path length of gas at 10 to 100 mm pressure usually gives a reasonable spectrum. I n the condensed phases, sample thickness usually must be less than 0.1 mm. Solutions are convenient but, of course, the solvent must be selected with attention to its own infrared spectrum. The two most commonly used solvents are CCl, and CSz, selected because of their relative transparency. Unfortunately water absorbs heavily through much of the infrared region and only extremely narrow "windows" are available for spectral studies.
The limits of applicability as a qualitative analytical tool are mainly fixed by the complexity of the sample. The more molecular types present, the more difficult it is to unravel the spectrum. In our fingerprint analogy, we can picture the difficulty of discerning a particular print in a pattern of several partially overlaid fingerprints. The same problem can be a limitation in quantitative use, superimposed on the ever-present accuracy limit fixed by the photometry. In general, it is reasonably easy to make quantitative determinations on simple mixtures to an accuracy of 5 to 10yo. With extreme care (and experience) it is possible to achieve accuracies in the range 0.5 to 1%. Figure 3 presents an impressive example of the "fingerprint" aspect of the vibrational absorptions. These are the spectra of the three trimethylbenzenes. Since the boiling points of these compounds are close together (data shown on Fig. 3) and the chemistries are pract,ically identical, normal analytical methods are relatively impotent. Yet there is not the slightest difficulty in differentiating the three compounds by infrared spectra or in analyzing mixtures of them.
What Are the Chemical Applications o f Infrared Spectroscopy?
Figure 3. Infrared spestro of the trimethylbenzene.; print" spectral region.
Infrared spectroscopy has such wide use in chemistry that it is difficult to list comprehensively the types of applications. There are, however, six uses w h i ~ hare of part,icular importance.
Figure 4 shows a more dramatic comparison. These are the spectra of two steroid hydrocarbons, as presented by Jones and Sandorfy (1). The two hydrocarbons, cholestane and coprostane, are diastereoisomers, differing mly in the spatial configuration about the juncture of the A and B rings, as shown in Figure 5. This conformational difference has profound effect in the infrared spectra, making identification of the pure substances a simple matter. The ability to distinguish such similar molecules gives the user of this technique great analytical power. Of course the method depends upon accessibility of reference spectra of the pure compounds. This need has been answered by vast compilations of spectra. The American Petroleum Institute has assembled over
Quolitotive ond Quontitotive Analysis
Each molecule has its own characteristic set of ahsorption frequenciesas distinctive as a fingerprint. The possibilities for qualitative detection of the presence of a given molecular species in a sample are obvious. In addition, the amount of the light absorbed by a sample is a quantitative index of the number of molecules absorbing. Thus, an infrared spectrum has the capability of answering both of the questions "What is in the sample?" and "How much is in the sample?"
WAVELENGTH
(MICRONS the
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Consideration of the chemistry of each of these products suggests the staggering analytical problem that would have been posed if conventional chemical methods were used. Diagnosis of Functional Groups
Figure 4. Infrared spectra of two rteroidr (From Jones, R. N., m d Sandorfy, C., "Technique of Organic Chemistry:' vol. IX, Chemical Applieotion3 of Spectroscopy. Interscience Publishers, Inc.. copyright 1956.)
2000 spectra of pure hydrocarbons (%). R. N. ,Jones and co-workers have presented spectra of over 300 steroids (3). Cannon and Sutherland have placed in the literature infrared spectra of 104 aromatics, including many carcinogenic compounds (4). Other extensive and "open-ended" compilations are growing continually (5, 6). The analytical use of infrared spectroscopy in a practical, albeit exciting, application is demonstrated by a recent research publication of Brown and Pimentel (7). These workers decomposed nitromethane, CH3N02,in a solid sample at low temperature by photolysis with ultraviolet light. The experimental conditions are sufficiently unusual that it was difficult to anticipate what products would be obtained. From a series of infrared spectra, however, it was possible to identify eight different products: methyl nitrite, formaldehyde, carbon monoxide, carbon dioxide, cyanic acid, nitrous oxide, nitric oxide, and water, not one of which had been expected. Neither of the two expected products, ethane and nitrogen dioxide, was formed. Not only were the eight products identified, but also the relative amounts could be deduced, providing a basis for a detailed scheme of the sequence of events which caused the formation of this complex mixture. All of this information was obtained from samples containing only a few micromoles of each product.
H GHOLESTANE
-
GOPROSTANE
Figure 5. The spotid sonformdions of the steroids cholestane and Eopratane.
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Many chemically important functional groups have characteristic infrared absorptions. This makes the infrared spectrum a useful tool in the study of a molecule of unknown structure. After a conventional elemental analysis has indicated which atoms are present, the infrared absorptions may show how they are bonded together. Some examples are presented in Tables 1 and 2. Table 1 shows the absorption regions characteristic of &H, N-H, G H , and S-H groups. Clearly one can distinguish the types of atoms to which hydrogen is attached. In the case of carbon-hydrogen linkages, it is even possible to differentiate acetylenes, olefins, and Table 1.
Characteristic Absorptions of Some Hydrogen Linkages
Functional eroun
- .
0-H N-H C-H
3650-3590 3500-3300 3300-3250 (acetylenes) 304%3010 (olefins) 2960-2850 (paraffins) 2600-2800
S-H
Table 2.
Frequency range in which absorption occurs (em-li
Characteristic Absorptions of Some Carbon Linkages
Functional eroun u
.
C=N C=C C=O C=C C-0 C-F C-CI c-Rr
Frequency range in which absorption occurs icm-') 2260-2215 2260-2100 1825-1650 1680-1620 1270-1060 1400-1000 800-600
fiw~nn
paraffins. Table 2 shows the characteristic ranges for some carbon linkages. These tables present the rosy side of the picture since a more complete listing would reveal overlaps which detract from the individuality. Nevertheless the capability of the method is great, particularly with the aid of Bellamy's systematic collection of characteristic spectral features (8). The absorption of the stretching motion of the C=O functional group has probably received more study than any other infrared feature. Extensive tables of carbonyl frequencies permit segregation of the functional groups which contain C=O and even identification of the molecular environment of the group. Bellamy brings much of this together in his book (8) and Table 3 shows the potentialities by reference to spectral ranges of some ketonic absorptions. References to the original literature and to other compilations are liberally given by Bellamy. An historic example of the use of infrared spectra in diagnosis of functional groups concerns the elncidation of the structure of penicillin. The infrared studies of Rasmussen and Brattain (9) significantly accelerated
the clarification of the structure of this important antibiotic. These workers, in a pioneering application of spectroscopy to complex molecules, were able to eliminate two incorrect, hut a t the time highly favored, structures, one a tricyclic structure and the other containing the oxazolone group. The latter structure, for example, was contra-indicated by the absence of the characteristic absorptions of the oxazolone group near 1820 cm-I (5.50 p) and -1670 cm-' (-6.0 fi). The presence of the monosubstituted amide linkage in the correct structure, shown in Figure 6, was deduced from the presence of an intense absorption near 1515 cm-' (6.6 p). More crucial, however, was the proposal of the (3-lactam functional group. This idea was based on the interpretation of the absorption of penicillin at 1780 cm-' (5.62 ,J) as the characteristic absorption of 8-lactams (1739 em-', 5.75 fi) shifted by fusion to the thiazolidine ring. Synthesis of suitable fused thiazolidine-p-lactam model compounds and examination of their spectra confumed the last interpretation. The work of Brown and Pimentel, already referred to (7), gives a more recent example. After the identification of the eight familiar substances listed, there remained distinct absorptions of another species of unknown structure. The sample was too small for any type of chemical analysis and this new material was identified by a diagnosis of the functional groups indicated by the frequencies recorded. The molecule nitroxyl, HNO, was thus detected for the fist time-identified by the presence of absorptions characteristic of the N-H group and the N=O group. Quonfifofive Deferminofiono f Molecular Strucfure
There are two ways in which infrared data yield detailed information about molecular structures (i.e., about bond lengths, hond angles, and molecular symmetry). If the spectmm of a pure substance in the gaseous state can be recorded, fine structure snperimposed on the vibrational absorptions can be studied. This fine structure (caused by molecular rotation) can be resolved readily for the smaller molecules. It has furnished information about the moments of inertia, hence the structures, of many simple molecules. Even for the larger molecules, however, there is a possibility of determining the amount of molecular symmetry by contrasting the infrared spectrum to spectral data of another type, Raman scattering data. A research problem in which both of these approaches were useful is the study of the structure of pentaborane, B6Hg,by Hrostowski and Pimentel (10). At the time this study was initiated, the structure of this unusual compound was not known. Coupling the Raman and infrared spectral data showed that the molecule possesses high symmetry. The contours of some of the bands led to estimates of the molecular moments of inertia. These and other facets of the spectra all were Table 3.
substituted amide group Figure 6.
thiazolldine ring
The structure of penicillin-G.
interpreted in terms of a four-sided pyramidal boron skeleton, the same structure deduced by other workers using X-ray and electron diffraction methods. Nitrous acid provides another interesting rase. Jones, Badger, and Moore (11) examined the infrared spectrum of the gaseous material as shown in part in Figure 7. The line structure, identified in the figure by the regularly spaced lines, permitted estimates of the molecular moments of inertia. This study provides today the best available information concerning the structure of this molecule, shown in Figure 8. These workers verified the existence of cis- and trans- forms, they deduced the atomic arrangement, and they estimated the bond angles and bond lengths of each form. I t has not yet been possible to measure these quantities by any other method of structure determination. Exominotion o f Chemicol Bonding
Tables 1 and 2 have already revealed the sensitivity of the infrared spectrum of a molecule to the exact nature of its chemical bonding. The different ahsorption frequencies of acetylenic, olefinic, and paraffinic C-H bonds reflect changes in the chemical bond between carbon and hydrogen as a result of the molecular environment. In Table 2, the higher frequency shown for C=O compared to C-0 honds is caused by the greater strength of the double bond. The same difference is shown for the carbon-carbon triple hond compared to the carbon-carbon double bond. The information concerning chemical hond strength can he extracted both through empirical correlations and through detailed mathematical analysis of the spectrum. The latter approach is called "normal coordinate analysis" and it yields, finally, numerical
The C = O Stretching M o d e of Ketones ( 8 )
Ketone type u-halogen Saturated openchain Aryl org unsaturated Diarvl
-4Br--CO -CH2-CO-CH2-Ph-CO-CH=CH-CO-Ph-CO-Ph-
Frequency range (om-') 1745-1725 1725-1705 17OCb1680 1685-1665 1670-1660
-J [cm-'1 Figure 7. Rotational Rne structure in the infrared spectrum of goseour nitrous acid (the regularly spaced lines identify features caused by rotation) ( 1 11.
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CIS
Figure 8.
The structures of cis- ond lranr-nitrous mid.
values for the "force constants" which restrain the molecule in its equilibrium configuration. These force constants based on a simple Hooke's Law approximation, correlate with chemically important properties, such as hond lengths and hond dissociation energy. We can return once again to the study of the photolysis of nitromethane (7) for an example. Having detected for the first time the nitroxyl molecule, Brown and Pimentel examined the nature of its chemical bonds. Calculation of force constants permitted characteriz:ttion of the hond orders in the molecule. Correlation of force constants and hond lengths of more well-understood molecules, as shown in Figure 9, proved that t,he nitrogen and oxygen atoms are linked by a double, not a single hond. This eliminates the possible structure H-0-N, in favor of H-N=O. Correlations of force constants with bond lengths such as that presented in Figure 9 have long been recognized for diatomic molecules. Recent extensions of such correlations (vibrational frequencies or force constants correlated with bond lengths) to bonds in polyatomic molecules include N-N bonds (IS), C-N bonds (It), and C-0 bonds (IS, 14, 15). Chemical effects, too, have been correlated with frequency shifts. Bellamy (8) summarizes attempts t,o find evidence of inductive effects, mesomerism, conjugation, and dipolar field effects in vibrational frequencies.
cal mechanics permit calculation of its thermodynamic properties to temperature ranges far beyond the limits of experimental measurements. It is appropriate to cite examples from the research of Pitzer and his coworkers. Their extensive calculations of the contributions of vihrational degrees of freedom to the thermodynamic properties of hydrocarbons are incorporated in the American Petroleum Institute Tables (16). Pitzer's work extends inclusively from ethane (17) and branched paraffins (18) through cyclohexane (19) to the substituted naphthalenes (20). The calculation of thermodynamic properties is frequently preceded by formidable problems of spectroscopic anaylsis. This statement inevitably brings to the writer's mind his investigation, with McClellan, of the infrared spectrum of naphthalene for the purpose of determining its thermodynamic properties (dl). The reference cited (21) concluded two years of research and a series of articles on more detailed aspects of the problem, made necessary by the complexity of the vibrational spectrum of this large molecule. With 18 atoms, naphthalene has 48 distinct vihrational frequencies, all of which had to be estimated reliably! With a complete vihrational assignment, McClellan and Pimentel calculated the temperature dependencies of the naphthalene heat capacity, enthalpy, free energy, and entropy from 300 to 1500°K. Study o f W e a k Chemical Inferoctions; The Hydrogen Bond
There are some important, though weak, chemical interactions which influence the infrared spectrum. The prime example is the hydrogen hond, an interaction in which a hydrogen atom already bonded strongly to one atom bonds weakly to a second atom. This type of bond is much weaker than a normal covalent hond, yet it affects noticeably many physical and chemical properties. Dramatic examples are provided by the unique properties of water (e.g., its high melting and boiling points) and by the structures of proteins. Proteins have characteristic helical configurations whose stabilities depend specifically upon hydrogen bonding. In view of the importance of hydrogen bonding, it is fortunate that this interaction causes significant and easily detectable disturbances of the infrared spectrum.
p
N204 NOi HNO p ClNO
-0
Colculotion o f Thermodynamic Properties
Chemical equilibrium is expressed quantitatively in terms of thermodynamic functions: energy, enthalpy, free energy, and entropy. The temperature dependencies of these functions are influenced by vihrational degrees of freedom. Hence a complete understanding of the vibrational spectra of molecules is essential to the prediction of the shift of equilibrium caused by a t,emperature change. Once the rharacteristic vihrational frequencies of a molecule are determined, then the methods of statisti656 / lourno1 o f Chemical Education
1.0 Figure 9.
1.2 R
a
Force conrlont versus bond length for N-0
1.4 bonds (71.
In fact, the spectral changes which accompany hydrogen bond formation are so readily observed and so unique in character that the infrared method is the easiest and probably most used method for studying hydrogen bonding. The characteristic stretching vibration of a hydrogen atom involved in a hydrogen bond is found a t lower frequency, with much greater absorption coefficient, and with unusual band width. The magnitudes of the changes are shown in Figure 10 for methanol. When it is dissolved in CCI, a t high dilution, no hydrogen bonds are formed. The spectrum under these conditions is the uppermost one in the figure. The rather narrow absorption near 3650 cm-' is caused by the 0-H stretching motion. I n the pure liquid and in the solid phase this absorption is altered drastically, revealing the extensive hydrogen bonding. The potentialities as a diagnostic tool are obvious. There are particularly important applications in the infrared detection of intramolecular hydrogen bonding. When a molecule possesses both an acidic functional group (such as an 0-H group) and a basic group, there is a possibility that a hydrogen bond will form within the molecule, protided the spatial arrangement is suitable. The presence or absence of such bonding is, then, a basis for determining the spatial relations among functional groups. To see the sensitivity of the method, consider the systematic work of Kuhn (82) with the cyclohexane diols. Some of the data collected by Kuhn are presented in Table 4. The last column Table 4.
0-H
Compound
Stretching Frequencies (cm-'1 of Diols
Fa(free)
didgin bonded)
A
i.
shows that the various conformations are readily distinguished through accurate frequency measurements. This conformational criterion has proven to be of great value in studies of more complex cyclic and polycyclic compounds. The most convincing testimonial to the value of the infrared study of hydrogen bonding is the immense volume of literature which has accumulated. The applications are discussed at length (and an extensive bibliography is given) by Pimentel and McClellan (25). There are cited there correlations of vibrational frequency shifts with a variety of chemically interesting properties: heat of hydrogen bond formation, heat of solution, hydrogen bond length., base strength, and Henry's Law constants. Summary
In summary, we have examined six types of applications of infrared spectroscopy to problems of great interest to the chemist. These uses span such a broad area of chemistry that we are left without wonder that infrared spectroscopy has become an everyday lab-bench tool. Recognizing the importance of the technique, we must be sure that it receives due emphasis in the undergraduate curriculum. Every chemist should know the potentialities of infrared spectral study and every research chemist
CH,OH In CCL
z50C
very dilute
crystal I
1 3600
3400
3200
s5,cni*
Figure 10. Infrared spectra of methanol in condensed phase* ~pectrolregion d t h e 0-H stretching motion 1231.
the
should consider its possible use in the solution of his own chemical problems (24). Litemlure Cited
( I ) JONES, R. N., A N D SANDORFY, C., "Tcehniqu~of Organic Chemistry," vol. IX, Chemical Applications of Spertroscopy, Interscience Publishers, Inc., Kcw York, 1956. (2) American Petroleum Institute Research Projrct 44, Infrared Spectral Data, Carnegie Institute of Teohnology, Pittsburgh, Pa. (3) DOBRINER, K., KATZENELLEXBOGEN, E , R., A N D JONES, R. N., "Infrared Absorption Spectra. of Steroids, an Atlas," Interscience Publislirrs, Inc., Y e w York, 1953. G . B. B. If.,Specbo(4) CANNON, C. G., AND SUTHERLAKD, chim. A d a , 4, 373-95 (1951). (5) National Research Council-Sationnl Bursau of Standards Keysort Cards (E. C. Creitz), Kstional Research Council Committee on Infrared Spectra, National Bureau of Standards, Washington 25, D. C. A N D SANDORFY tabulate other compilations; see ref(6) JONES erence 1, pp. 328-29. H. W.. A N D PIMENTEL. G. C.. J . C ~ C V PI .~"v s..29. . . (7) BROWN, 883-88 (1958). (8) . . BELLAMY.L. J.. "The Infrared S ~ e e t r aof Comr,lvz hlolccules,"'2nd ed., John Wiley & s&, Ine., New $&, 1958. (9) "The Chemistry of Penicillin," Princeton Univ. Presr, Princeton, New Jersey, 1949. H. J., A N D PIMENTEL, G. C., J . Ant. C h m . (10) HROSTO~SKI, Sac., 76,098 (1954). L. H., BADGER, R. M., A N D MOORE,G. E., J . Chem. (11) JONES, Phys., 19,1599 (1951). W. J., Chem. Reus., 57, 1179-1211 (1957). (12) ORVILLE-THOMAS, V. A,, J . (13) LAYTON,E. M., KROSS, R. D., A N D FASSEL, Chem. Phys., 25,135 (1956). M., FILLWALK, F., FASSEL,V. A,, A N D RUN(14) MARGOSHES, DLE, R. E.,J. ChemPhys., 22,381 (1954). W. E.: A N D PIMENTEL, G. C., J . (15) EWING,G. E., THOMPSON, Chem. Phys., 32,927-32 (1960). (16) Ross1~1,F. D., ET AL., "Selected Values of Physical and Thermodynamic Properties of Hydrocarbons and Related Comnounds." Csrneeic Press. Pittsbureh. - , 1953. (17) KEMP,5. D., 'AND PI&ER, K.'s., J. Am. Chcm. Soc., 59, 278-79 (1037). (18) PITZER,K. S., Chem. Reus.,27,30-57(1940). (19) BECKETT,C. W., PITZER,K . S., A N D SPITZER,R., J . Am. Chem. Soc., 69,248 (1047). E. D., A N D PITZER,K. S., J . Ant. (20) MILLIGAN, D. E., BECKER, Chem. Soc., 78,2707 (1956). G. C., J . Chem. Phys., (21) MCCLELLAN, A. L., AKD PIMENTEL, 23,24548 (1955). (22) K u n ~L. , P., J. Am. Chem. Soc., 74,2492-99 (1952). G. C., A N D MCCLELLAN, A. L., "The Hydrogcn (23) PIMENTEL, Bond," W. H. Freeman & Co., San Francisoo, 1960. (24) Reference 1 is an excellent source of further information concerning the use of infrared spectroscopy in chemistry. Volume 37, Number 12, December 1960
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