6672
J . Phys. Chem. 1989,93, 6612-6615
Gas-Phase Structures of the Hydrogen Fluorophosphoranes PHFl and PH2F, Dines Christen," Johannes Kadel,Ia A. Liedtke,IbRolf Minkwitz,lb and Heinz Oberhammer*JP Institut fur Physikalische und Theoretische Chemie, Universitat Tiibingen, 7400 Tiibingen, West Germany, and Institut fur Anorganische Chemie, Universitat Dortmund, 4600 Dortmund, West Germany (Received: February 21, 1989)
The geometric structures of the phosphoranes PHF, and PH2F3have been determined by joint analyses of electron diffraction intensities and rotational constants. The rotational constants of PHF, have been reported in the literature; those of PH2F3 have been derived from the microwave spectrum in the present study. The harmonic vibrational corrections for interatomic distances and rotational constants were calculated from general harmonic force fields based on the vibrational frequencies. The following geometric parameters (rrvalues with 3u error limits) were obtained. PHF,: PF, = 1.539 (3)A, PF, = 1.596 (3) A,PH = 1.380 (15) A, F,PF, = 114.5 (5)O, F,PF, = 90.6 (3)'. PH2F3: PF, = 1.539 (5) A,PF, = 1.618 (4)A,PH = 1.398 (15) A, HPF, = 117.1 (17)O, F,PF, = 91.9 (4)O.
Introduction Conformations, geometric structures, bonding properties, and dynamic behavior of pentacoordinated phosphorus compounds have attracted great interest over the past decades. Excellent reviews of these topics have been given by R. R. Holmes2 and by R. L ~ c k e n b a c h . ~Acyclic phsophoranes in general adopt nearly ideal trigonal-bipyramidal structures (tbp), and large deviations from this ideal tbp structure have been observed only in cyclic or polycyclic compounds. In monodentate phosphoranes, the distribution of various ligands to axial and equatorial positons and the pseudorotational exchange of these ligands have been of major interest. The ligand site preference is to a large extent governed by electronegativity and by Electronegative ligands favor axial positions, whereas *-donating ligands prefer equatorial positions. In some compounds, the conformation is affected by steric interactions as well.lo Furthermore, structures and conformations of pentacoordinated phosphorus compounds are important as transition states or metastable intermediates in chemical reactions where their conformations determine the stereochemistry of the products." Various models have been proposed for the bonding in these hypervalent phosphorus compounds, which differ with respect to d-orbital participation. Pauling, Nyholm et al., and Gillespie et al. rationalized the formation of five bonds by full incorporation of one d orbital in the formation of sp3d hybrids,I2-l4 whereas Rundle et al. suggested an electron-deficient model with threecenter four-electron bonds in the axial direction, which does not require d-orbital p a r t i c i p a t i ~ n . Ab ~ ~ initio calculations indicate the latter bonding model, but d functions on phosphorus are required to reproduce experimental structures correctly, and
( I ) (a) Universitat Tiibingen. (b) Universitat Dortmund. (2) Holmes, R. R. Pentacoordinated Phosphorus; ACS Monographs 175 and 176; American Chemical Society: Washington, DC, 1980; Vol. I and 11. (3) Luckenbach, R. Dynamic Stereochemistry of Pentaco-ordinated Phosphorus and Related Elements; Thieme: Stuttgart, FRG, 1973. (4) Muetterties. E. L.; Mahler, W.; Schmutzler, R. Inorg. Chem. 1963, 2, 613. (5) Gillespie, R. Molecular Geometry; Van Nostrand: London, 1972. (6) Hoffmann, R.; Howell, J. M.; Muetterties, E. L. J . Am. Chem. SOC. 1972, 94, 3047. (7) McDowell, R. S.; Streitwieser, A,, Jr. J. Am. Chem. SOC.1985, 107, 5849. (8) Strich, A,; Veillard, A. J. Am. Chem. SOC.1973, 95, 5574. (9) Kutzelnigg, W. Angew. Chem. 1984, 96, 262, and references cited therein. (10) Oberhammer, H.; Grobe, J.; Le Van, D. Inorg. Chem. 1982.21, 275. ( I 1) Mislow, K. Am. Chem. Res. 1970, 3, 321. (12) Pauling, L. The Nature of the Chemical Bond, 3rd ed.; Cornell University Press: Ithaca, NY, 1960. (13) Craig, D. P.; Maccoll, A,; Nyholm, R. S.; Orgel, L. E.; Sutton, L. E. J. Chem. SOC.1954, 332. (14) Gillespie, R. J.; Nyholm, R. S. Q.Rev., Chem. SOC.1957, 1 1 , 339. ( 1 5 ) Rundle, R. E. J. Am. Chem. SOC.1963, 85, 112.
0022-3654/89/2093-6672$01.50/0
despite a small contribution, d functions contribute strongly to the stability of these compounds. In addition to the covalent bonding, polar contributions may play an important role as well. The hydrogen fluorine phosphorane series PH,Fs,, n = 0-5, have frequently been used as model compounds in the discussion of bonding properties, and a number of theoretial investigations have been reported in the l i t e r a t ~ r e . ~As ~ our ~ ~ experimental '~~~ studies were being completed, we learned of a high-quality ab initio study on the geometries and vibrational spectra of the entire seriesz8 The existence of PHF, and PH2F3was reported about 20 years ago by R. R. Holmes,29P. M. T r e i ~ h e land , ~ ~B. Blaser et aL31 Their conformations were determined by vibrational and N M R spectra, which were interpreted in terms of the equatorial position of the hydrogens, Le., C , symmetry for both compounds. Pierce and C o r n ~ a lderived l ~ ~ rotational constants for PHF4 and its deuterated species from the microwave spectra. These experimental data, however, were not sufficient for a complete structure determination. So far, no structural information has been reported for PH2F3. PH3F2is a rather unstable species, and it has been characterized by vibrational and N M R spectra.33 PH4F and PH5 have not yet been observed experimentally. According to a b initio calculations,8 the barrier to pseudorotation increases from PFS (4.8 kcal/mol) to PHF, (9.2 kcal/mol) to PH2F3(1 3.0 kcal/mol). The latter result is in agreement with the value estimated by Holmes (10-15 kcal/m01).~~Thus, both molecules can safely by treated as rigid rotors, performing small amplitude vibrations. In this study, we report experimental gas-phase structure determinations for PHFl and PH2F3. Because no further stable Kutzelnigg, W.; Wasilewski, J. J . Am. Chem. Soc. 1982, 104, 953. Keil, F.; Kutzelnigg, W. J. Am. Chem. SOC.1975, 97, 3623. Walker, W. J. Mol. Spectrosc. 1972, 43, 411. Rauk, A.; Allen, L. C.; Mislow, K. J . Am. Chem. Soc. 1972,94,3035. Altmann, J. A.; Yates, K.; Csizmadia, I. G. J. Am. Chem. SOC.1976, 98,'1450. (21) Shih, S.-K.; Peyerimhoff, S. D.; Buenker, R. J. J . Chem. SOC.,Faraday Trans. 2 1979, 75, 379. (22) Kutzelnigg, W.; Wallmeier, H.; Wasilewski, J. Theor. Chim. Acta 1979. 51..~ 261. (23) Trinquier, G.; Daudey, J.-P.; Caruana, G.; Madaule, Y. J. Am. Chem. SOC.1984. 106, 4794. (24) Reed, A. E.; Schleyer, P. v. R. Chem. Phys. Lett. 1987, 133, 553. (25) Bestmann, H. J.; Chandrasekhar, J.; Downey, W. G.; Schleyer, P. v. R. J. Chem. SOC.,Chem. Commun. 1980, 978. (26) Strich, A. Inorg. Chem. 1978, 17, 942. (27) Marsden, C. J. J . Chem. SOC.,Chem. Commun. 1984, 401. (28) Breidung, J.; Thiel, W.; Komornicki, A. J. Phys. Chem. 1988, 92, 5603. (29) Holmes, R. R.; Storey, R. N. Inorg. Chem. 1966, 5 , 2146. (30) Treichel, P. M.;Goodrich, R. A.; Pierce, S. B. J . Am. Chem. SOC. 1967, 89, 2017. (31) Blaser, B.; Worms, K-H. Z . Anorg. Allg. Chem. 1968, 361, 15. (32) Pierce, S.B.; Cornwell, C. D. J . Chem. Phys. 1968, 48, 21 18. (33) S e d , F.; Velleman, K. Z . Anorg. Allg. Chem. 1971, 385, 123. (34) Holmes, R. R. Acc. Chem. Res. 1972, 5 , 296. 1
-
0 1989 American Chemical Society
Gas-Phase Structures of the Hydrogen Fluorophosphoranes
n
0.0
5.0
The Journal of Physical Chemistry, Vol. 93, No. 18, 1989 6673 TABLE I: Exwrimental R o h t i o ~ Transitions l (MHz) of HIPFI
K-1'
J' 1
10.0
,
I
15.0
1
I
25.0
20.0
30.0
1
35.0
l/Rngstrom
s in
Figure 1. Experimental (dots) and calculated (full line) molecular intensities and differences for PHF,.
t7
I
I
I
I
I
0
5
10
15
20
---..
25
K+I'
J" 0
K-1" 0 1 0 2 1 2 1 2 1 3 1 4 0 4 1 5 1 5 0 6 1 7 4 7 3 7 1 8 4 8 3 4 6 9 5 8 1 1 6 7 1 1 6 9 1 3 7 8 1 3 7 IO 5 16 9 1 5 8 10 15 8 IO 6 17 6 18 11 11 17 9 10 17 9 11 18 10 10 18 IO 12 7 20 1 2 3 0 1 1 3 2 3 4 4 5 7 5 7 8
1 1 0 2 2 2 1 2 2 1 2 3 2 2 3 2
2 3 2 2 3 4 4 5 5 6 8 8 7 9 9 1 0 1 2 5 12 5 14 6 14 6 15 11 16 7 16 7 16 11 17 12 18 8 18 8 19 9 19 9 19 13
K+I" 0 1 2 1 2 2 4 3 4 5 5 4 4 6 4 5 5 5 6 6 7 6 8 7 7 7 8 9 8 9 8
Vexp
- Vcalc -0.02 0.04 0.09 0.02 0.06 -0.01 0.11 -0.01 0.02 -0.06 -0.04 0.02 0.03 0.01 0.04 -0.04 -0.05 0.01 0.03 -0.08 -0.01 -0.02 -0.09 -0.04 0.04 0.1 1 0.02 0.04 0.01 0.06 -0.11
Vexp
15855.78 23 195.70 19 104.16 21 890.25 25 547.42 20683.99 15 946.55 19 717.25 19608.00 20 882.47 20 533.20 19 194.47 21 202.08 22815.33 22631.36 24 274.21 19719.10 20 221.42 20944.56 22 339.89 22 590.50 25 156.78 24 421 .08 24 338.25 16 118.95 23 252.87 26 300.27 26 326.85 18 799.96 18 803.23 21 335.72
TABLE 11: Rotational Constants (MHz) and Centrifugal Distortion Constants (kHz) of H2PFI
I d
30
35
Figure 2. Experimental (dots) and calculated (full line) molecular intensities and differences for PHzF3.
B'dexpt) 12 185.88 (15) 4939.95 (6) 3 669.94 (6)
A B C Dj
0.82 (0.07)
= 6.59 (0.48) DK = 11.8 (0.2) DjK
fluorine and phosphorus isotopes exist, rotational spectroscopy alone does not allow for complete and accurate structure determinations for these compounds. On the other hand, the radial distribution functions (see text below) demonstrate that gas electron diffraction by itself does not provide enough structural information either for a reliable determination of all geometric parameters. Thus, only joint analyses of microwave and electron diffraction data lead to complete and accurate structures for these phosphoranes. The rotational constants of PHF, and its deuterated species were taken from the literature;32 those of PH2F3were derived from its microwave spectrum. The harmonic vibrational corrections were calculated from harmonic force fields based on the vibrational frequencies. So far, we were not successful in synthesizing PH3F2 with the purity required for an electron diffraction study. Because of its vanishing static dipole moment (DUsymmetry), the microwave spectrum cannot be recorded with conventional techniques.
Experiments PHF4 and PH2F3were synthesized according to methods reported by Holmes et al.29 that are based on the reaction of phosphoric and hypophosphoric acids, respectively, with HF. The purity of the samples was checked by 'H, I9F, and 31PN M R spectra (Bruker Model AM 300) and was better than 99% for both compounds. The gas IR spectra were recorded with a Perkin-Elmer Model 580B spectrometer and were found to agree with those reported by Homles et with two exceptions in the case of PHF,: the absorptions at 1528 and 629 cm-I, which are (35) Holmes, R. R.; Hora, C . J., Jr. Inorg. Chem. 1972, 11, 2506.
B'Jexpt) 12 161.0 (4) 4932.0 (12) 3 668.4 (3)
B',(calc)' 12 161.3 4933.0 3 668.7
6 j = 0.21 (0.04) 6K
= 3.22 (0.67)
'Calculated with r, parameters of Table VI
TABLE 111: Diagonal Valence Force Constants (N/cm) for PH,F+, ( n = 0-2)'
APFJ jlPF,) jlPW AFePFe) AHPFJ flHPH) AFePF,) AHPF,)
PFS
PHF4
PHP,
6.47 5.45
6.18 5.23 3.59 0.41 0.18
5.95 4.80 3.69
0.28 2.24
2.00 1.07
0.25 0.16 1.72 0.96
"Complete force fields can be obtained from the authors upon request.
reported with medium intensities, are very weak ( HPF,. All trends in bond angles are predicted correctly by the ab initio calculations, although the calculated deviations of the axialequatorial angles from 90' are smaller than the experimental values. Acknowledgment. Financial support by the Fonds der Chemischen Industrie is gratefully acknowledged. Furthermore, we thank Prof. W. Thiel for sending us his manuscript prior to publication. Registry No. PHF,, 13659-66-0; PHzF3, 13659-65-9. (45) Maki, A. G.; S a m , R. L.; Olson, W. B. J . Chem. Phys. 1973, 58, 4502. (46) Morino, Y . ;Kuchitsu, K.; Moritani, T. Inorg. Chem. 1969, 8, 867.
