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Studies on the Chemistry of Halogens and of Polyhalides. XXVI. Infrared Absorption Spectra of Iodine Bromide and Its Complexes
by Yoshiharu Yagi, Alexander I. Popov, Department of Chemistry, Michigan State Univereitu, Eaet Lansing, Michigan 48823
and Willis B. Person Department of Chemistry, Univeraity of Florida, Gaineaville, Florida $2601
(Received October 17, 1966)
The influence of complexation on the frequency of the I-Br fundamental stretching vibration in iodine bromide complexes has been investigated. The frequency and the intensity of this band are strongly influenced by the strength of the iodine bromide-electron donor interaction. The position of the IBr band varies from 261 cm-I in carbon tetrachloride solutions to 205 cm-’ for the pyridine-IBr complex dissolved in benzene. I n solid complexes the IBr band is split into a doublet. The changes in frequency are analyzed to obtain changes in force constants. Some attempt is made to interpret these results.
Introduction Ultraviolet and visible spectroscopic studies of iodine and of iodine monochloride complexes with various organic electron donors have been abundantly reported in the literature.’ A much smaller amount of work has been done on the infrared spectra of such complexes. Several studies have been reported on the infrared spectra of halogen-aromatic compound complexes in which the authors investigated the changes in the spectra of the donor molecules produced by the formation of the donor-acceptor complex.2-e The influence of complexation on the fundamental stretching frequency of iodine monochloride has been described in one of the previous publications of this series.’ It was found that the I-CI stretching frequency is quite sensitive to the strength of interaction between the interhalogen and the donor molecule with which it formed a complex. The frequency decreased steadily from 375 cm-I for the uncomplexed iodine monochloride in carbon tetrachloride solution to 275 cm-1 for the strongest complex of the series, pyridine-IC1. At the same time, the intensity of the band gradually increased with the increasing strength of the complex. Similar behavior was reported for iodine complexes.* Although iodine stretch could not be observed in non-
complexing solvents such as carbon tetrachloride or cyclohexane, since iodine, obviously, does not have a dipole moment, polarization of the 1-1 bond by complex formation renders the vibration infrared active. The N-I stretching frequency in Iz-trimethylamine complexes has been reported by Yada, et al.,Qto be at 145 cm-’ and infrared absorption bands for the N-I and I-X stretching vibrations have been recently measured for a series of iodine chloride and iodine bromide complexes with pyridine and several substituted pyridines. loJ (1) L. J. Andrews and R. M. Keefer, “Molecular Complexes in Organic Chemistry,” Holden-Day, Inc., San Francisco, Calif., 1964, and G. Briegleb, “Elektronen-Donator-Acceptor-Komplexe,” Springer-Verlag, Berlin, 1961. (2) W. Haller, G. Jura, and G. C. Pimentel, J . Chem. Phys., 22, 720 (1954). (3) T. Collin and L. D’Or, ibid., 23, 397 (1955). (4) D. L. Glusker and H. W. Thompson, J. Chem. Soc., 471 (1955). (5) E. E. Ferguson, J . C h m . Phy8., 25, 577 (1956). (6) M. Marzocchi and E.Ferroni, Gazz. Chin. Ital., 91, 1200, 1216 (1961). (7) W. B. Person, R. E.Humphrey, W. A. Deskin, and A. I. Popov, J . Am. Chem. SOC.,80, 2049 (1958). (8) M. V. Lorenzelli, Compt. Rend., 258, 5386 (1964). (9) H. Yada, J. Tanaka, and S. Nagakura, J . Mol. Spectry., 9 , 461 (1982). (10) A. G. Maki and R. D. Nelson, submitted for publication.
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Y. YAGI,A. I. POPOV, AND W. B. PERSON
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Table I : Source and Purity of the Solvents Used in This Investigation Source
Purification
Not purified Dried over sodium and fractionally distilled
Carbon tetrachloride Cyclohexane
Eastman Spectrograde Matheson Coleman and Bell, reagent grade Aldrich Chemicals, reagent grade Matheson Coleman and Bell Matheson Coleman and Bell
1,ZDichloroethane
Matheson Coleman and Bell
1,ZDibromoethane
Aldrich Chemicals, reagent grade Matheson Coleman and Bell
Solvent
Acetonitrile Ben I ene Carbon disulfide
n-Heptane Nitrobenzene Nitromethane Pyridine Toluene &Xylene
Matheson Coleman and Bell, reagent grade Matheson Coleman and Bell Eastman, reagent grade Eastman, reagent grade Aldrich Chemicals, reagent grade
Not purified Purified by previously described techniquelo Dried over granulated BaO and fractionally distilled Washed with sulfuric acid, water, and dilute solution of sodium carbonate in that order. Refluxed over P2O5and distilled twice from the same reagent Distilled Dried over granulated BaO and fractionally distilled Dried over BaO and distilled Dried over Pz05and fractionally distilled Dried with BaO for 24 hr and fractionally distilled Not purified Not purified
It is interesting to note that while numerous complexes of iodine and of iodine monochloride have been investigated and their formation constants have been determined, the same generality does not apply to the complexes of iodine bromide. In fact, with the exception of mesitylene, l2 pyridine and substituted pyridines,lOzlsand some nitrile^,'^-'^ the strengths of iodine bromide-donor interactions seem to be unknown. This study, therefore, was initiated in order to determine more completely the complexation ability of iodine bromide and the influence of such complexation on the fundamental frequency of the IBr molecule.
obtained. The product was filtered, washed with carbon tetrachloride, and dried in vacuo. It has a decomposition point of 95-looo. Anal. Calcd for ClaHsNIBrzIt: C, 21.08; H, 1.42; N, 4.91; iodometric equivalent, 142.5. Found: C, 20.95; H, 1.45; N, 4.85; iodometric equivalent, 143.0. Several attempts were made to prepare iodine bromide complexes of 2,2‘-biquinoline and 1,10-phenanthroline by similar techniques, but they were unsuccessful. Spectroscopic Measurements. Infrared absorption
Experimental Section Reagents. Iodine bromide was prepared and puri-
(11) 5. G.W. Ginn and J. L. Wood, Trans. Faraday SOC.,62, 777 (1966). (12) J. H. Blake and R. M. Keefer, J . Am. Chem. Soc., 77, 3707 (1955). (13) A. I. Popov and R. H. Rygg, ibid., 79, 4622 (1957). (14) A. I. Popov and W. A. Deskin, ibid., 80, 2976 (1958). (15) W. B. Person, W. C. Golton, and A. I. Popov, ibid., 85, 891 (1963). (16) P. Klaboe, ibid., 84, 3458 (1962). (17) P. Klaboe, ibid., 85, 871 (1963). (18) P. Klaboe, Acta Chim. Scand., 17, 1179 (1963). (19) E.Augdahl and P. Klaboe, ibid., 19, 807 (1965). (20) J. Cornog and R. A. Karges, Inorg. Syn., 1, 165 (1939). (21) Y.A. Fialkov and I. D. Muzyka, Zh. Obshch. Khim., 18, 1205 (1948). (22) A. I. Popov, J. C. Marshall, F. B. Stute, and W. B. Person, J . Am. Chem. SOC.,83, 3586 (1961).
fied by the procedure of Cornog and Karges.20 The 2,2’-bipyridine was obtained from the G. F. Smith Chemical Go. and was purified by recrystallization and the 4,4’-bipyridine was a gift from Professor J. C. Marshall, St. OIaf College. The source and purity of the solvents used in this investigation are given in Table I. The preparation of pyridine-IBr and 4,4’-bipyridine21Br complexes has been described in the literature.21*22 A new complex, 2,2’-bipyridine-2IBr, was prepared by slowly mixing carbon tetrachloride solutions ( ~ 0 . M 1 ) of the two components. A fine yellow precipitate was The Journal of Physical Chemistry
STUDIES ON
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CHEMISTRY OF HALOGENS AND
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spectra were obtained with a Perkin-Elmer Model 301 spectrometer. Solid samples of the complexes were dispersed in Nujol. The mulls were sandwiched between polyethylene disks. Solution measurements were carried out in polyethylene cells of 0.5-, 1.0-, and 4.0-mm path length. The concentration of the solutions studied ranged from 0.03 to 0.25 M in IBr. While iodine bromide solutions are somewhat less , reactive than those of iodine chloride and, consequently, 100 IEQ 200 2x, fewer experimental difficulties were encountered than WAVENUMBER, cm" in the previous study17nevertheless the spectra usually Figure 1. Infrared absorption spectra of the showed some change after the solutions were aged for pyridine-halogen complexes: -, Py-IC1 (Nujol) ; - - - -, Py-IBr (Nujol); . . ', Py-IBr (benzene solution). several hours. Consequently, all measurements were made on freshly prepared solutions. Polyethylene cells have a tendency to swell when left in contact with organic solvents for any length of time. This, in turn, changes the optical path length and makes it very difficult to evaluate accurately and reproducibly the absorption intensities. Since l,Bdichloroethane, nitrobenzene, and o-xylene have strong absorption bands in the 300-200-~m-~ region, the donor solvent was diluted with carbon tetrachloride. It was found that 0.1-0.2 M IBr dissolved in a 90-10% v/v mixture of carbon tetrachlorideIC4 IM 200 250 300 WAVENUMBER, om? donor solvent gave entirely satisfactory spectra. I
Results and Discussion The general features of the infrared absorption of the iodine bromide complexes are shown in Figures 1 and 2. Absorption spectra of two iodine chloride complexes are included for comparison. It should be noted that the I-Br stretching vibration near 195 cm-l, the I-CI stretch near 270 cm-l, and the N-I stretch in the IC1 complexes near 175 cm-' all split into doublets in the spectra of the solid complexes. This splitting has also been observed by Maki and Nelson,lo who have attributed it to correlation splitting due to more than one molecule in the unit cell. Careful comparison of our spectrum of the Py-IC1 complex with that of Maki and Nelson'O shows that the relative intensities of the two bands near 175 cm-l are reversed for the two samples. This observation is consistent with the assignment of this pair as a correlation-field doublet, together with the suggestion that some (differing) amount of orientation and polarization by the spectrometer optics existed in the two samples. Further support for the assignment of this pair is provided by the observation that the spectrum of Py-IBr dissolved in benzene shows only a single symmetrical peak near 200 cm-l. It is of some interest to note, however, that the N-I stretching vibration near 159 cm-l for the solid Py-IBr complex does not split, in contrast with the behavior for the Py-IC1
t
Figure 2. Infrared absorption spectra of the 4,4'-bipyridine-halogen complexes (Nujol): , 4,4'-BP-BICI; - - -, 4,4'-BP-2IBr.
-
complex where both the I-C1 and N-I stretching vibrations split by about the same amounts. The frequencies for the solid complexes are reported in Table 11. In Table 111, we list frequencies found for IBr dissolved in a series of solvents. Our values in Table I1 and Table I11 agree with those found by Maki and Nelsonlo for the two pyridine complexes and with those found by Stammreich, Forneris, and T a v a r e ~ ~ ~ and by Ginn and Wood11 within the accuracy of our frequency measurement. (There is an apparent calibration error, so that our results are uniformly 3 cm-' lower than those found by Maki and Nelson.'O) The fundamental absorption frequency for IBr in the gaseous state is 267 cm-1 (268.4 cm-' for 1127Br79 and 266.4 cm-1 for I127Bfll). The shift from this value to that (261 cm-') found in the noncomplexing solvents (CCL, n-heptane, and cyclohexane) is similar to the shift observed previously for IC1 (from 382 cm-' in the gas phase to 375 cm-1). It is apparently an "ordinary" solvent shift due to the charge in the effective field;24 the additional shift to lower frequencies shown (23) H.Stammreich, R.Forneris, and Y. Tavares, Spectrochim. Acta, 17, 1173 (1961).
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Table I1 : Halogen Vibrations in Solid IBr and IC1 Complexes
Complex
Pyridine-IBr PyridineIC1 4,4’-Bipyridine2IBr 4,4’-Bipyridine2ICl 2,2 ’-Bipyridine-2IBr 2,2 ‘-Bipyridine -2IC1
I-x
N-I
stretch, cm-1
stretch, cm-1
199, 190 274, 266 194.5, 187 274 211, 203 284.5, 268
158.5 178, 169.5 ( s ) 138 154.5 135 151 (s), 130.5 (w)
Table 111: Fundamental Vibration Band of IBr in Various Solvents
Solvent
Carbon tetrachloride n-Heptane Cyclohexane 1,Z-Dichloroethane” Nitrobenzene” Nitromethane Carbon disulfide 1,ZDibromoethane Benzene Toluene +Xylene” Acetonitrile Pyridineb
I-Br band, om-’
Ak/krBr A k / k I c i
Formation constant of the IBr complex in CClr a t 2 6 O , 1. mole-]
261.0 261’ 261.0 261.0 259.0 254.5 253.0 253.0 249.0
0.05 0.06 0.06
249.0 251’ 247.0 247.0 244.5 205.0 204’
0.09 0.10 0.10 0.12 0.38
0.07
0.10
=0.2c
0.13
11.4d 1.4 x 1 0 4 ~
The measurements were made on IBr dissolved in a mixture of donor solvent-carbon tetrachloride (10-90%). Solid Py-IBr complex was dissolved in benzene. Estimated from comparing the formation constant of mesitylene-IBr with mesitylene-It, mesitylene-IC1 and of benzene-Is and benzene -1C1, L. J. Andrews Referand R. M. Keefer, J. Am. Chem. SOC.,74,4500 (1952). ence 14. ‘ Reference 13. Reference 23. ’ Reference 11.
‘
’
in Table I11 is an indication of the sacrificial chargetransfer interaction.'^^^*^^ As usualj7there is a borderline region (the solutions in dichloroethane, nitrobenzene, nitromethane, and carbon disulfide) where it is not clear whether the shift represents “ordinary” solvent shift or charge-transfer interaction,27 but then the solutions in stronger donors show shifts to lower frequencies as the strength of the donor increases, just as observed for the IC1 c o m p l e x e ~ . ~ * ~ ~ On examining Tables I1 and 111, we note that the shift in frequency observed for the I-Br stretching The Journal of Physical Chemietry
vibration is less than that observed for the I-C1 stretching vibration in complexes with the same donor. This observation might suggest that the interaction between IBr and these donors is weaker than that between IC1 and these same donors. A similar conclusion might be reached on comparison of the N-I stretching frequencies for the two complexes, since those for the IBr complexes are about 20 cm-l lower than those for the IC1 complexes. However, we must be quite cautious about drawing such conclusions from the frequency data alone. The reason for caution is that the magnitude of sacrificial donor-acceptor action is measured, not by the frequency shift, but rather by the change in the I-X force constant. For many complexes the I-X stretching frequency is isolated from the higher frequency vibrations of the donor and from the much lower D-A stretching vibrations. Hence, the shift in frequency is just proportional to the change in force constant and either one can be used to measure the extent of donoracceptor action. However, for strong complexes, the D-A stretching vibration (D-A = N-I for these complexes) increases in frequency until it is comparable to that for the I-X stretching vibrations. When that happens, the two vibrations are expected to mix and split apart. As this interaction continues in a series of stronger and stronger complexes, eventually the highfrequency vibration changes its nature from pure I-X stretch to pure N-I stretch and vice versa for the lowfrequency stretch. In the frequency region where the two vibrations mix strongly, it will be necessary to do a normal coordinate analysis to find the force constants. It is entirely possible for one complex to have a lower ‘“-1 stretching frequency” and a higher “I-X stretching frequency’’ than another complex and still have X larger ~ N I indicating , a stronger complex smaller ~ I and but greater mixing of the vibrations. Since the I-Br stretching vibration already comes a t lower frequency in the free molecule than does the I-Cl stretching vibration, we must expect this mixing to occur for weaker complexes than will be the case for the IC1 complexes. Hence, we must attempt to (24) A. D. Buckingham, Proc. Roy. Soc. (London), A248, 169 (1958);A255,32 (1960). See also the discussion by H. Yamada and W. B. Person, J. Chem. Phys., 40, 309 (1964). (25) R. S. Mulliken and W. B. Person, Ann. Rev. Phys. Chem., 13, 107 (1962),and references cited therein. (26) H. B. Friedrich and W. B. Person, J . Chem. Phys., 44, 2161 (1966). (27) We believe, however, that the shift is large enough to indicate
interaction. (28) It should be noted that the failure of Maki and Nelson to observe significant trends is presumably because all the donors they studied show a comparable interaction. See the discussion below.
STUDIES ON THE CHEMISTRY OF HALOGENS AND
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POLYHALIDES
calculate the force constants (kmr and ~ N Ifor ) the IBr complexes. For the weak complexes V N I is expected to be very low (less than 100 cm-l) ; hence, we may use the force constants computed for a diatomic molecule, as we have done before.’ For the strong aminehalogen complexes, we must use equations for a linear triatomic molecule, as done by Maki and Nelson.lo We may use these equations without apology because the other vibrations of the molecule are expected to be separated so far that no interaction with these two frequencies should occur. The results of these calculations are presented in Table IV.
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tical for these amine-halogen complexes, while the values of kIx vary considerably when the possibility for mixing with the D-I stretching vibration is considered under the “triatomic molecule approximation.” If we assume that the D-I bond increases in strength as the weight of the dative structure (b2) increases, in a treatment paralleling that for the I X force conthen we may compute the value of the D-I force constant in the pure dative structure, ~D-I’. For the amine-IX complexes, we should expect this constant to be the same for all complexes. From the three complexes for which both the N-I and the I-X stretching vibrations are identified, we see that the value of ~ N I ’ is indeed remarkably constant-2.5 0.05 mdynes/A. We may test further the validity of this treatment by using the data for the trihalide ions.1° This test is presented in the lower part of Table IV. We see there that the predicted force constants for the D I bonds are lower than the known values by about 0.5 mdyne/A. Hence, the value for the NI force constant ( ~ N I ’ = 2.5 mdynes/A) may be too low by about 0.5 mdyne/A. If so, then the frequency for the N-I stretch in the Py1 2 complex may be about 120 cm-l. Ginn and Wood“ report absorption at that frequency for this system, which they attribute to the N-I stretching vibration. However, Yarwood and Person have been unable to reproduce their results,31so there may be some question. At any rate the force constant ~ C N I O = 2.5 mdynes/A should be accurate enough to predict approximate positions for the N-I stretching vibration in other complexes. The N I force constant, ~ N I ,listed in Table IV for the (CH3)3N-12 complex is different from that reported by Yada, Tanaka, and Nagakura;8 we were not able to reproduce their calculated value, but found the value listed in Table IV instead. Finally, we note that the values for the force constants listed in Table IV do depend on our assumption that klz = 0.40 mdyne/A. However, the dependence of these force constants (kIx and kNI) on klz is not very great.1° The ideas presented here are believed to be almost quantitatively correct, but the uncertainties in klz and in VIX and V N I due to the abnormally large solvent effect” result in uncertainties of the order of &O.l unit in force constants. One point of confusion in the interpretation of the force constants of complexes should now be discussed.
*
Table IV: Force Constants (mdynes/A) for Some Halogen Complexes (D-IX) Complex
kIxa
(CH3)3N-I2 Py-IC1 Py-IBr Py-Le
1.21 1.24 1.27 1.21
I --I?
c1--IC1° Br --IBrh
kixb
Ak/kbrc
0.97 0.45 1.36 0.42 1.27 0.36 ( 1 . 3 5 ) (0.23) 0.92 0.47 1.02 0.57 0.91 0.56
kDIb
1.15 1.02 0.88 (0.58) 0.53 1.02 0.91
kDI
2.55 2.44 2.45 (2.50) 1.1 (cf. 1 . 7 ) 1.8(cf.2.4) 1.6(cf. 2.1)
’
From From diatomic molecule approximation (see text). t,riatomic molecule approximation (see text). Defined by ( k ~ x ”- k~x)/krx”. ~ D I ”is the force constant for the D-I single bond in the pure dative structure D+-IX; the value listed here is comput,ed from ~ D I ’ = k ~ ~ / ( A k / k )e . For the values listed under the triatomic molecule approximation for this compound, we have assumed that the N-I stretching vibration comes a t 115 cm-I. Values computed for CsI3. See ref 10. O Computed for (CH3)4NIC12. See ref 10. * Computed for (CH3)&1Br2. See ref 10.
’
As Maki and Nelson have noted,’O there are three unknown force constants for the linear triatomic molecule: k N I , ~ I X ,and the interaction constant, klz. For strong complexes, the latter is expected to be abnormally large, as discussed by RIaki and Nelson,1° based on the same reasoning they have been applying to the trihalide ions.29~30 Maki and Nelson have evaluated force constants for amine-halogen complexes assuming different values of klz, with klz = 0.40 as the preferred choice.1° We agree with their analysis and we present in Table IV the results based on that X under “dichoice for klz. The value of ~ I listed atomic molecule approximation” represents one extreme value for this conetant and the value under “triatomic molecule approximation” represents the other extreme. We note in Table IV that the “diatomic molecule approximation” gives results for k ~ which x are all iden-
(29) W. B. Person, G. R. Anderson, J. N. Fordenwalt, H. Stammreich, and R. Forneris, J . Chem. Phgs., 3 5 , 908 (1961). (30) A. G. Maki and R. Forneris, submitted for publication. (31) J. Yarwood and W. B. Person, submitted for publication.
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We have noted the expectation that the k’s will change progressively as the “strength of the complex’’ increases. To measure the latter, one often uses the formation constant of the complex. Actually, the relates A k / k to the weight of the dative structure (b2) in the wave function of the complex, \ k ~ . The strength (measured by Kr) of the complex is also related to b2, if that strength is determined primarily by the chargetransfer interaction. However, it is quite certain that other factors also contribute to the magnitude of Kr. For example, IC1 complexes might be expected to be more stable than the corresponding I 2 complexes because of the contribution to the stability of the former by dipole-induced-dipole forces. Actually, Table IV suggests that the charge-transfer contribu-
tion is greater for IC1 complexes than for Izcomplexes. However, one should remember that the force constants and the formation constants do measure different things, so the correlation between them may not always be very good. Hence, for example, the fact that A k / k (Table 111)for the weaker IBr complexes (with benzene and other weak donors) is almost the same as that found for the corresponding IC1 complexes does not seem especially surprising. Acknowledgment. This work was supported by the
U. S. Army Research Office, Durham. The authors are also indebted to Dr. A. G. Naki of the National Bureau of Standards for a preprint of his papers and to Professor John C. Marshall, St. Olaf College, for the 4,4‘-bipyridine.
The Corroding Iron Surface. I. Dissolution of Iron in the Halogen Acids
by E. McCafferty and A. C. Zettlemoyer Center for Surface and Coatings Research, Lehigh University, Bethlehem, Pennsylvania (Received October 18, 1966)
The corrosion rates for iron in deaerated 1 N solutions of the halogen acids at 25” were determined by cathodic polarization and substantiated by analytical methods to be 42, 23, and 10 pa/cm2 for HC1, HBr, and HI, respectively. The observed cathodic Tafel slopes were also well ordered: HC1 < HBr < HI. Anodic Tafel slopes calculated from the Stearn-Geary equation, on the other hand, were not markedly dependent on the nature of the anion. The observed order of corrosion rates may be explained on the basis that the desorption of the halide anion at anodic sites is the rate-determining step in the corrosion process.
Introduction The study of corrosion is of interest to several disciplines, each approaching the subject with a different viewpoint. The purpose of these two papers is to investigate the corrosion properties of a well-defined system and to interpret the observations in terms of principles. The system chosen was iron in solutions of the halogen Mkb at 25”. This article presents the experimental methods and results. The Journal of Physical Chemistry
In the second paper, the nature of the surface during corrosion will be discussed in terms of these results.
vi^^ of Theory The concept of mixed ~otentialsl-~and the measurement of corrosion rates by the polarization method (1) C.Wagner and W. Traud, 2.Elektrochem., 44, 391 (1938). (2) I. Kolthoff and C. S. Miller, J. Am. Chem. Soc., 62, 2171 (1940). (3) J. v. Petrooelli, J . Electrochem. soc., 97, io (1950).
