Chemical Education Today
Letters
Infrared Spectroscopy and Bond
Strengths
A recent article in this Journal (1) emphasizes the
relationship between infrared frequencies, bond lengths,
and bond strengths. Nevertheless, it contains two significant errors that can be used to illustrate the difficulties in using infrared data to get information about
bond strengths.
The first problem found in the interpretation of infrared spectra is band assignment. Indeed, the authors
assign to S–S stretching vibrations one band at 1242 cm–1
in the infrared spectrum of K2S2O6 and one band at 1635
cm –1 that they found in the spectrum of (hydrated?)
Na2S2O3 (1). However, both bands appear at a frequency
too high to be assigned to a S–S stretching vibration,
because for S=SF2 and gaseous S2 molecules, which contain a double S=S bond, υS=S appears at 718 cm –1 (2).
In fact, for Na2S 2O3 υS–S has been reported at 446 cm,
while for K2S2O6?2H2O one polarized Raman band at 293
cm–1 has been assigned to υS–S (the dithionate anion is
centrosymmetric and S–S is not active in the infrared
spectrum (2, 3). Actually, the band at 1242 cm–1 in the
infrared spectrum of K2S2O6 should be assigned to an S–O
stretching vibration (2, 3). On the other hand, the band
at 1635 cm–1 that the authors assign to υS–S in Na2S2O3
(1) is at a frequency too high to be assigned even to S–O,
because these bands appear at 1123 and 995 cm–1 in the
vibrational spectra of Na2S2O3 (2). A likely explanation
is that the authors isolated hydrated Na2S2O3, and that
the band at 1635 cm–1 corresponds to a deformation of
the water molecules.
Once the bands have been properly assigned, another difficulty in correlating band positions with bond
strength comes from coupling. Indeed, bond strengths are
related to the force constants, but band positions may
not be directly related to the force constants if the vibration of interest is strongly coupled to other vibrations.
A striking example of this phenomenon is provided by
the infrared spectra of ionic cyanate and thiocyanate
salts. The principal resonance structures of the cyanate
and thiocyanate ions are show in Figure 1. Bearing in
mind that the C–N stretching frequencies are 2168 cm{1
for OCN– and 2049 cm–1 for SCN–, the authors concluded
that the C–N bond is stronger, and hence the participation of resonance structures I is higher, for cyanate than
for thiocyanate ions (1). Nevertheless, all the available
data indicate that the opposite situation is found. Indeed,
the NC force constant is higher in NCS– than in NCO–
(4, 5). This fact suggest that the weight of resonance
structure I is higher for thiocyanate than for cyanate,
and in agreement with this idea the negative charge on
the N atom is higher for OCN– than for SCN– (4,5). Recent calculations indicate that the atomic charges in
NCO– are –0.64 on N and –0.38 on O, while in NCS– they
are –0.46 on N and –0.54 on (6). Furthermore, the CN
bond order is higher and the EC (E = O,S) bond order is
lower in SCN– than in OCN – (4, 5). Therefore, it is clear
that resonance structure I is more important in SCN–
than in OCN–, and that the CN bond is stronger in SCN–.
In agreement with this conclusion, while the C–N distance in ionic thiocyanate salts is close to 1.17 Å (7), the
calculated C–N distance in NCO– is 1.19 Å from both ab
A296
initio studies (8) and matrix infrared spectra of condensed phase NCO– (9).
In order to understand why υCN appears at a higher
frequency in NCO– than in NCS –, although the CN bond
is stronger in the latter ion, we must take into account
the υCN and υCE (E=O,S) are not pure vibrations, but
they are coupled to each other giving rise to an antisymmetric NCE stretching (mainly υCN) and a symmetric
NCE stretching (mainly υCE) (4). Bearing in mind that
υCO appears at a higher frequency than υCS, its frequency is closer to that corresponding to CN, so that coupling is more important in NCO– than in NCS–, thus raising the frequency of the antisymmetric NCO stretching
vibration υ(CN). Indeed, recent calculations of the atomic
displacements during the antisymmetric stretch show
that while this mode is significantly delocalized over the
molecule in NCO–, it is essentially a CN stretching mode
in NCS– (6). The case of NCO – and NCS– provides a very
good example, for teaching purposes, of the misconclusions
that can be obtained in correlating vibrational frequency
with bond strength. Indeed, even the widely accepted idea
that shortening of a bond reflects strengthening has been
recently challenged (10).
According to the accepted rules for drawing Lewis
structures, a structure is considered stable if negative
formal charges are assigned to atoms with higher electronegativity (11). Therefore, it may be surprising that
resonance structure I below has more weight for SCN–,
while resonance structure II has more weight for OCN–,
although the negative formal charge is not placed on the
more electronegative atoms:
E
C
E
N
I
C
N
II
E = O, S
The higher weight of resonance structure II for OCN– than
for SCN– may be related to the higher tendency of second
row elements to use p orbitals in p bonding as compared
to third and subsequent row elements (12).
Literature Cited
1. Wiskamp, V.; Fichtner, W.; Kramb, V.; Nintschew, A.; Schnieder, J. S. J. Chem.
Educ. 1995, 72, 952.
2. Weidlein, J; Muller, U.; Dehnicke, K. Schwingungsfrequenzen I; Georg Thieme:
Stuttgart, 1981; pp 93, 94, 147, 148.
3. Palmer, W. G. J. Chem. Soc. 1961, 1552.
4. Norburn, A. H. Adv. Inorg. Chem. Radiochem. 1975, 17, 231.
5. Golub, A. M.; Kohler, H.; Skopenko, V. V. Chemistry of Pseudo-Halides;
Elsevier: Amsterdam, 1986; Chapter 1.
6. Li, M. ; Owrutsky, J.; Sarisky, M; Culver, J. P.; Yodh, A; Hochstrasser, R. M. J.
Chem. Phys. 1993, 98, 5499.
7. Wells, A. F. Structural Inorganic Chemistry, 5th ed.; Clarendon: Oxford, 1986;
p 935.
8. Cai, Z. L. Chem. Phys. 1993, 170, 33.
9. Smith, D. F, Jr.; Overend, J.; Decius, J. C.; Gordon, D. J. J. Chem. Phys. 1973,
58, 1636.
10. Ernst, R. D.; Freeman, J. W.; Stahl, L; Wilson, D. R; Arif, A. M.; Nuber, B;
Ziegler, M. L. J. Am. Chem. Soc. 1995, 117, 5075.
11. Ahmad, W. Y.; Omar, S. J. Chem. Educ. 1992, 69, 791.
12. Huheey, J. E.; Keiter, E. A.; Keiter, R. L. Inorganic Chemistry, 4th ed.; Harper
Collins: New York, 1993; p 861.
David Tudela
Departamento de Quimica Inorganica
Universidad Autonoma de Madrid
28049–Madrid, Spain
Journal of Chemical Education • Vol. 73 No. 12 December 1996