PACIFIC SOUTHWEST ASSOCIATION O F CHEMISTRY TEACHERS THE EVIDENCE FROM INFRARED SPECTROSCOPY FOR HYDROGEN BONDING
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A Case History of the Correlation and Interpretation of Data MEL GORMAN University of San Francisco, San Francisco, California
T H E nature and phenomena associated with the hydrogen bond, together with some of the experimental methods used to investigate it, have been discussed in THIS JOURNAL (1). This paper deals with infrared spectroscopy, which in the last twenty years has proved t o he one of the most powerful tools for the study of hydrogen bonding. Most compounds exhibit very complicated absorption when subjected to the whole range of the infrared spectrum. This absorption reflects the rotation of molecules and all of the various modes of interatomic vibration within the molecule which cause bonds to be deformed by stretching, bending, twisting, etc. Each of these types of rotatory and oscillatory motion absorbs its characteristic amount of radiant energy, and a large number of them may be detected a t various frequencies in the spectrum of a given compound. Fortunately, both from the theoretical and exoerimental standooint. some of the most important and useful correlations in the study of hydrogen bonding have been derived from a study of the stretching frequency of the X-H bond as observed in the near infrared. For a detaileddiscussion of the theory and application of infrared absorption in hydrogen bonding the sources in reference (3) may be consulted. A brief but excellent treatment of experimental technique is given by Ewing (3) and a very complete presentation, including a l i t of catalogues and atlases of spectra, is given by Jones and Sandorfy (4).
EFFECT O F HYDROGEN BONDING ON INFRARED ABSORPTION
The effect of hydrogen bonding by infrared absorption can be illustrated by considering the 0 - H . .0 bond as an example. If a compound such as an alcohol is in the gaseous state, its molecules on the average are so far apart that there can be no formation of R H
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dimers, R 4 , . , H-0, through hydrogen bonding. Thus the "free," i.e., non-hydrogen bonded, 0-H groups can vibrate by stretching a t an unperturbed frequency. These oscillations are comparatively in-
dependent of the structure of the rest of the compound since the hydrogen has such a small mass. Thus it is found that the hydrogenic stretching of the free 0-H group in almost all gaseous compounds will absorb infrared radiation sharply a t a frequency close to 3700 cm.-'. However, in the liquid or solid states or in relatively concentrated solutions, association through hydrogen bonding to form dimers and higher polymers takes place. The O--H group is no longer "free." Its absorption is shifted to lower frequencies by hundreds of wave numbers in many compounds and flattened to give a broad band. Since the change of frequency of a few wave numbers is easily measured, it is obvious that the method is a sensitive one for detecting hydrogen bonds. The lowering of the frequency is to be expected, since the strength of the originally free 0-H bond is weakened when it becomes nart of the 0-H . .0 bond. COLLECTION O F DATA AND EMPIRICAL CORRELATION
Interpreting the phenomena of hydrogen bonding constitutes an interesting case history illustrating the development of chemical science. 4 study of almost any topic in chemistry wbich has had a long enough history will reveal that its unfolding can be divided into three stages. First there is the gathering of the data. As soon as enough information is available the second stage of empirical correlation appears. Finally the time will be ripe for a theoretical interpretation, often with the modification of a simple original theory which may have been used as a framework for the first two stages. Of course, these categories do not appear in neat chronological sequence, for there is much overlapping, especially between the first two. At the present time, hydrogen bonding presents an instance in which all three stages are evident. Since the modern concept of the hydrogen bond originated only about thirty-five years ago (6), it should not be surprising that there is still much gathering of experimental data. Although the use of infrared spectroscopy in the study of hydrogen bonding has become widespread only in the last twenty years, it is a very JOURNAL OF CHEMICAL E ~ C A T I O N
popular technique because it can be applied to substances in the solid, liquid, gaseous, or dissolved state, and because in the last decade there have been important strides in instrumentation. Moreover, it is a powerful tool because it can detect weak bonding and chelation. From the theoretical point of view, the method is attractive because it gives a direct measure of the forces involved. I n this respect it differs from other useful methods such as X-ray diffraction, which can locate the hydrogens only by inference, and entropy studies, in which calculations must be made from heat capacity studies. One of the most important pioneer developments in the application of infrared absorption to hydrogen bonding is the work of Hilbert, Wulf, Hendricks and Liddell in 1936 (8). This contribution and those of other earlier investigators are reviewed by Pauling (7). The interesting point t o note in the paper of Hilbert and his co-workers is that they speak of the absence of the characteristic (free) absorption bands in hydrogen bonded substances. However, in 1937 Badger and Bauer (8)pointed out that in a strict sense these bands do not disappear, but are shifted and broadened, and that any decisions concerning hydrogen bond formation must be based on the shifts and general character of the absorption. I n a second paper in the same year Badger and Bauer (9) continued their study of this subject, and attempted a correlation of the stretching frequency shift, A T , (v for free 0-H absorption, minus v for the hydrogen bonded 0-H), with the strength of the hydrogen bond. Unfortunately, spectroscopic data for only six compounds were available whose energies of hydrogen bond formation were known with fair accuracy. Nevertheless a semiquantitative linear relationship was found when Av was plotted against energy. I n pursuing this problem still further with additional data, Badger (10) concluded that the relationship is not quite linear as earlier supposed. He also emphasizes the rough nature of the correlation, due to errors in both the energy and spectral data. Although later considerations (11) made it evident that frequency shifts could not be associated indiscriminately with proportionality of hydrogen bond strength, the work of Badger is important for two reasons. In the first place, it has established the importance of the stretching frequency shift, which later provided the basis of extensive correlation with interatomic distances between hydrogen bonded atoms, and secondly it provides an example of the emergence of the second phase of the historical development of its topic, using the bare minimum of available data. It should he emphasized that there is need for more data even at the present time. Such compounds as organic substances with hydrogen bonded water of crystallization, aliphatic and alicyclic amines with N-H. . .N bonds, and substances having hydrogen bonded chlorine should be investigated. By 1950 an extensive literature of X-ray research on the interatomic distances between hydrogen bonded atoms had accumulated, but much of the work was not accurate enough for discussion of the bond. Then Donohue (1%)undertook a critical review of the data and selected values of the interatomic distance which had an accuracy of 0.1 A. or better. On the basis of Donohue's evaluation, Rundle and Parasol (18) were VOLUME 34, NO. 6, JUNE, 1957
able to choose eleven compounds which fell on a smooth curve in a plot of interatomic distance against the stretching frequency. Below 2.7 A. the absorption is very sensitive to differences in int,eratomic distance. This is the reason why only the most accurate X-ray data are suitable for correlation. Above 2.7 A the sensitivity falls off exponentially. This was followed by a somewhat similar treatment by Lord and Merrifield (14), who plotted Au against interatomic distance and found that eleven compounds under study fell on a straight line. The curve shows very strikingly that the shorter the bond the greater is the stretching frequency shift from the nou-hydrogen bonded value (3700 c m . 3 for 0-H and 3400 cm.-' for N-H). The year 1955 saw an important advance in the correlation between stretching frequency and interatomic distance in the paper of Nakamoto, Margoshes, and Rundle (16). This group extended the relationships of Rundle and Parasol to over 50 compounds containing eight types of hydrogen bonds X-H . .Y, in which X and Y may be fluorine, oxygen, nitrogen, and chlorine. I n a plot of interatomic distance against stretching frequency there are eight curves. The 0-H.. .O curve shows the same characteristics a.s that of Rundle and Parasol described above. Nakamoto, Margoshes, and Rundle presume that their other curves would show similar characteristics if more complete data were available. They also follow the procedure of Lord and Merrifield by plotting Av against interatomic distance for the 0-H . . . O bond, and get a similar linear relationship insofar as the data for the two groups of workers coincide, i.e., up to a bond length of about 2185 A. For longer bonds Av becomes smaller exponentially. Empirical correlation was extended still further by Pimentel and Sederholm (16),~vhopresented the relationship between interatomic distance and Av for 0-H. . .O, N-H . . .0, and N-H . . .N graphically and analytically in the form of an equation of the best straight line for each type, together with standard deviations in Au. THEORETICAL INTERPRETATION
Since the study of hydrogen bonding is still largely concerned with the task of acquiring more data, with a consequent limitation on empirical correlations, it is understandable that the third stage in historical development, theoretical interpretation, is as yet only in its infancy. There is a lack of full agreement as to the nature of the hydrogen bond, one school of thought adhering t o the simple and older electrostatic model, and another preferring a quantum mechanical formulation. Unfortunately, neither of these positions has been capable of offering a quantitative explanation of all of the features of the bond. But recently Lippincott and Schroeder (17) have proposed a very simple model for 0-H . . .O bonds in crystals, based on a potential function involving bond dissociation energy and interatomic distances, together with the assumptions of a van der Waals repulsion and an attractive electrostatic potential between the oxygen atoms. By appropriate operations with this function these workers are able to correlate and predict quantitatively a number of the main features of the hydrogen bond, such as the relationship between interatomic distance and hydrogen bond energy and between interatomic
distance and Av. Thus a direct relation between Av and bond energy can be obtained. Lippincott and Schroeder are careful to point out that their model is not intended as a complete one, for there are important aspects of infrared absorption which it does not explain. But they are optimistic that their model will be applicable to hydrogen bonding involving other atoms besides oxygen, and that it can be modified for liquid systems. Nevertheless, this limited model constitutes a definite advance from qualitative to quantitative theoretical interpretation. I t emphasizes the need of modifying the commonly held belief in an electrostatic model, in favor of quantum mechanical formulations, iaorder that our notions of the hydrogen bond can be improved. Lord and Merrifield ( I k ) , believe that an explanation of the spectroscopic properties of the hydrogen bond undoubtedly ill require a quantum mechanical model.
(8) BADGER, R. M., A N D S. H. BAUER,J. Chem. Phys., 5, 605608 (1937). R. M., AND S. H. BAUER,J. Chem. Phys., 5, 851 (9) BADGER,
R. M., J. Chem. Phys., 8,288 (1940). (10) BADGER, L. N., op. eit., pp. 260-61; M. M. DAVIES, (11) FERGUSON, Ann. Repts. Chem. Soe. (London), 43, 19 (1946). (12) DoNonuE, J., J. Phys. Chem., 56, 502-10 (1952). R. E., AND M. PARASOL, J. Chem. Phys., 20, 1487 (13) RUNDLE,
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(14) LORD,R. C., AND R. E. MERRIFIELD, J. Chem. Phys., 21, 166 (1953). K., M. MARGOSEES, A N D R. E. RUNDLE, J. (15) NAKAMOTO, Am. Chem. Soc.. 77.6480 (1955). (16) PIMENTEL, G. C., ~~ND'c. H. S E D E ~ O L M , J. Chem. Phys., 24, 639 fl!Xfi). J. Chem. Phys., 23, (17) L&N&TT,'E. R., AND R. SCHROEDER, 1099-1106 (1955).
LITERATURE CITED (1) JONES, E. V., J. CHEM.EDUC.,22, 76 (1945); R. W. TAFT, JR., AND H.'H. SISLER,ibid., 24, 175 (1947); M. GORMAN, ibid., 33, 468 (1956). L. N., "Electronic Structures of Organic Mole(2) FERGUSON, cules," Prentice-Hall, New York, 1952, pp. 229-36 and 251-68; L. KELLNER,Repts. Pmg. Phys., 15, 1-22 AND C. SA ORFY in "Technique or (1952); R. N. JONES Oreanie Chemistrv. " , Val,-IX. %mica1 Anolications of Spectroscopy," W. WEST, Editor, Interscience Publishers, New York, 1956, Chap. IV. (3) EWING,G. W., "Instrumental Methods of Chemical Analysis," McGraw-Hill Book Co., Inc., New York, 1954, pp.
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(4) J o s ~ s R. , N.,AND C. SANDORFY, op. eit., pp. 256-87 and 294-331. (51 Hr;ccrNs. RZ. L.. J. Am. Chem. Soe.. 53.3190 (1931): G. N.
p. 109. (6) HILBERT,G. E., 0.R. WULP,S. B. HENDRICKS, AND U. LIDDELL, J. Am. Chem. Soe., 58, 548-55 (1936). (7) PAPLING,L., "Nature of the Chemical Bond," Cornell University Press, Ithaca, 1940, pp. 316-27.
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