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Nitryl fluoride-xenon hexafluoride complex
3624 Notes J. inorg, nucl. Chem., 1973,Vol. 35, pp. 3624-3625. PergamonPre~s.Printed in Great Britain.
Nitryl fluoride-xenon hexafluoride complex (Received 15 December 1972) MANY metal fluoride and oxide fluorides react readily with nitrosyl fluoride and nitryl fluoride to form complexes[l-3]. Similarly, xenon hexafluoride and xenon oxide tetrafluoride combine with nitrosyl fluoride at room temperature[4] but with nitryl fluoride, although crystalline material is formed, the reactions do not go easily to completion. This, we find, also tends to happen when the hexafluoride is TcF 6 or ReF6[3 ]. On the other hand, MoF 6, WF 6 and U F 6 readily form NO2MF7 type complexes at room temperature[l] while PtF 6 reacts to give NO2PtF 6 and F 2 gas[5]. We describe here an investigation of the reaction between NO2F and XeF6, resulting in the formation of XeF6.NO2F.
EXPERIMENTAL A Monel and nickel vacuum manifold was used for all experimental manipulations. Reactions were carried out in thin-walled ( ~ ) in) nickel weighing cans fitted with small brass valves. Stoichiometries of the reactions were followed by weighing both starting materials, and reaction products in tared reaction vessels. The vacuum system and reaction vessels were passivated with fluorine and chlorine trifluoride prior to use.
Materials Xenon hexafluoride was prepared by the high pressure reaction of excess fluorine with xenon. The hexafluoride was purified from the other binary fluorides of xenon by taking advantage of the reversible reaction of XeF 6 with sodium fluoride. The detailed purification procedures are described elsewhere[6-8]. Nitryl fluoride was obtained from the Ozark-Mahoning Company. It was purified by pumping off non-condensible materials at - 1 9 5 ° and subsequent trap-to-trap distillation until its vapour pressure[9] and i.r. spectrum[10] were consistent with those cited in the literature.
X-ray diffraction measurements X-ray diffraction studies were carried out in well seasoned KeI-F capillaries.
In[tared and Raman spectra I.R. spectra were obtained with a Beckman IR-12 recording spectrophotometer. The Raman equipment, which has been described elsewhere[l 1] includes a Spectra-Physics 125 He-Ne laser and a Spex 1400 I1 double monochromator. I. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11.
J. R. Geichman, E. A. Smith and P. R. Ogle, Inorg. Chem. 1, 1012 (1963). N. Bartlett, S. P. Benton and N. K. Jha, Chem. Comm. 168 (1966). J. H. Holloway and H. Selig, J. inorg, nucl. Chem. 30, 473 (1968). G. J. Moody and H. Selig, Inorg. nucl. Chem. Letters 2, 319 (1966). F. P. Gortsema and R. H. Toeniskoetter, Inorg. Chem. 5, 1217 (1966). I. Sheft, T. M. Spittler and F. H. Martin, Science 145, 701 (1964). R. D. Peacock, H. Selig and I. Sheft, Proc. chem. Soc. Lond., 285 (1964). J. G. Maim, F. Schreiner and D. W. Osborne, Inorg. nud. Chem. Letters. 1, 97 (1965). O. Ruff, W. Menzel and W. Neumann, Z. anorg, allg. Chem. 208, 293 (1932). R. E. Dodd, J. A. Rolfe and L. A. Woodward, Trans. Faraday Soc. 52, 145 (1956). H. H. Claassen, H. Selig and J. Shamir, Appl. Spectroscop. 23, 8 (1969).
Notes
3625
RESULTS AND DISCUSSION The compound NOzF.XeF6 was prepared by the reaction of excess of NO2F with XeF6 at 100°. In a typical experiment XeFe (1.87 m-mole) in a nickel can was treated with a six-fold excess of NO2F for 4 hr. The reaction was quenched by rapid cooling with liquid nitrogen. There was no measurable pressure in the can at - 195°, indicating the absence of gaseous fluorine or oxygen. A pressure of about 400 mm was observed at - 7 8 °. This corresponds to the vapour pressure of NOzF at that temperature[9]. An i.r. spectrum of the volatile material showed only the absorptions associated with NO2F[10]. This excess NO2F was pumped off at - 7 8 ° to constant weight. The weight increase indicated an NO2F uptake of 1.75 m-mole corresponding to a stoichiometry 1.07 XeFe.NO2F. At room temperature the uitryl beptafluoroxenateO/I) is a colouriess crystalline solid with a vapour pressure of about 10 mm. The complex can be readily sublimed. In Kei-F tubes small rhombus shaped plates grow slowly over a period of several days. At 30° crystal growth is rapid and well formed crystals can be obtained in 24 hr. These crystals tend to be twinned and produce overgrowths after several days. Sometimes faults and cracks develop after subjection to X-rays. Powder diffraction patterns show similarities to NO2WF7 and NO2TcFT[3], but the product is not iso-structural with these compounds. I.R. spectra of the vapour in equilibrium with solid NO2F.XeF+ show only bands of the component molecules indicating considerable or complete dissociation in the vapour. In the solid deposited on cold AgC1 windows at - 195°, absorptions at 585 and 2300 cm-1 are observed. Absorptions in nitryl fluoride complexes in the 2360-2400 c m - t region indicate the presence of the NO2 + cation; those around 570600 cm - 1 are indicative of Xe-F stretching modes or of the bending mode of the NO 2 + cation. The Raman spectrum of the solid shows an intense sharp band at 636 cm- 1, medium intensity bands at 562 and 524 cm- 1 and a number of weaker hands in the bending region. On the other hand, a weak Raman hand at 2305 cm- ~ may possibly be due to the antisymmetric stretching vibration of the nitryl cation resulting from breakdown of the selection rules in the solid. The symmetrical stretching vibration of NO2 + which should be Raman active at about 1400 cm- t appears to be extremely weak. This, and the fact that the i.r. allowed mode is lower than the normally occurring 2380 cm- 1 suggest incomplete fluoride ion transfer. Moreover, the symmetrical Xe-F stretching vibration at 636 cm- 1 is not shifted substantially from the corresponding one at 657 cm-1 in solid XeF e, whereas the symmetrical Xe-F stretching vibration in XeFs" is shifted down considerably to 540 cm-114, 12]. The latter ion has been definitely established in the solid state[12]. The complex NO2F.XeFe can probably not be formulated as the ionic form NO2 +XeFT-, but rather as an intermediate fluorine bridged structure. The complex NO2F.XeF6 hydrolyses rapidly to form a colourless solution. Although some gas bubbles are evolved during hydrolysis, the resulting solution possesses oxidizing properties characteristic of aqueous xenon(VI). When solutions+exposed to the atmosphere dry out, the residual solid is explosive. Although analogous complexes of XeFe are known in the salts CsXeF7 and RbXeFT[13], the NO2F adduct is the only one of this type which is volatile and amenable to vacuum manipulations.
Chemistry Department The University oJ Leicester Leicester LEI 7RH Department o[ Inorganic and Analytical Chemistry The Hebrew University Jerusalem Israel
JOHN H. HOLLOWAY
HENRY SELIG URI EL-GAD
12. S. W. Peterson, J. H. Holloway, B. A. Coyle and J. M. Williams, Science 173, 1238 (1971). 13. R. D. Peacock, H. Selig and I. Sheft, J. inorg, nucl. Chem. 28, 2561 (1966).
Nitryl fluoride-xenon hexafluoride complex (Received 15 December 1972) MANY metal fluoride and oxide fluorides react readily with nitrosyl fluoride and nitryl fluoride to form complexes[l-3]. Similarly, xenon hexafluoride and xenon oxide tetrafluoride combine with nitrosyl fluoride at room temperature[4] but with nitryl fluoride, although crystalline material is formed, the reactions do not go easily to completion. This, we find, also tends to happen when the hexafluoride is TcF 6 or ReF6[3 ]. On the other hand, MoF 6, WF 6 and U F 6 readily form NO2MF7 type complexes at room temperature[l] while PtF 6 reacts to give NO2PtF 6 and F 2 gas[5]. We describe here an investigation of the reaction between NO2F and XeF6, resulting in the formation of XeF6.NO2F.
EXPERIMENTAL A Monel and nickel vacuum manifold was used for all experimental manipulations. Reactions were carried out in thin-walled ( ~ ) in) nickel weighing cans fitted with small brass valves. Stoichiometries of the reactions were followed by weighing both starting materials, and reaction products in tared reaction vessels. The vacuum system and reaction vessels were passivated with fluorine and chlorine trifluoride prior to use.
Materials Xenon hexafluoride was prepared by the high pressure reaction of excess fluorine with xenon. The hexafluoride was purified from the other binary fluorides of xenon by taking advantage of the reversible reaction of XeF 6 with sodium fluoride. The detailed purification procedures are described elsewhere[6-8]. Nitryl fluoride was obtained from the Ozark-Mahoning Company. It was purified by pumping off non-condensible materials at - 1 9 5 ° and subsequent trap-to-trap distillation until its vapour pressure[9] and i.r. spectrum[10] were consistent with those cited in the literature.
X-ray diffraction measurements X-ray diffraction studies were carried out in well seasoned KeI-F capillaries.
In[tared and Raman spectra I.R. spectra were obtained with a Beckman IR-12 recording spectrophotometer. The Raman equipment, which has been described elsewhere[l 1] includes a Spectra-Physics 125 He-Ne laser and a Spex 1400 I1 double monochromator. I. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11.
J. R. Geichman, E. A. Smith and P. R. Ogle, Inorg. Chem. 1, 1012 (1963). N. Bartlett, S. P. Benton and N. K. Jha, Chem. Comm. 168 (1966). J. H. Holloway and H. Selig, J. inorg, nucl. Chem. 30, 473 (1968). G. J. Moody and H. Selig, Inorg. nucl. Chem. Letters 2, 319 (1966). F. P. Gortsema and R. H. Toeniskoetter, Inorg. Chem. 5, 1217 (1966). I. Sheft, T. M. Spittler and F. H. Martin, Science 145, 701 (1964). R. D. Peacock, H. Selig and I. Sheft, Proc. chem. Soc. Lond., 285 (1964). J. G. Maim, F. Schreiner and D. W. Osborne, Inorg. nud. Chem. Letters. 1, 97 (1965). O. Ruff, W. Menzel and W. Neumann, Z. anorg, allg. Chem. 208, 293 (1932). R. E. Dodd, J. A. Rolfe and L. A. Woodward, Trans. Faraday Soc. 52, 145 (1956). H. H. Claassen, H. Selig and J. Shamir, Appl. Spectroscop. 23, 8 (1969).
Notes
3625
RESULTS AND DISCUSSION The compound NOzF.XeF6 was prepared by the reaction of excess of NO2F with XeF6 at 100°. In a typical experiment XeFe (1.87 m-mole) in a nickel can was treated with a six-fold excess of NO2F for 4 hr. The reaction was quenched by rapid cooling with liquid nitrogen. There was no measurable pressure in the can at - 195°, indicating the absence of gaseous fluorine or oxygen. A pressure of about 400 mm was observed at - 7 8 °. This corresponds to the vapour pressure of NOzF at that temperature[9]. An i.r. spectrum of the volatile material showed only the absorptions associated with NO2F[10]. This excess NO2F was pumped off at - 7 8 ° to constant weight. The weight increase indicated an NO2F uptake of 1.75 m-mole corresponding to a stoichiometry 1.07 XeFe.NO2F. At room temperature the uitryl beptafluoroxenateO/I) is a colouriess crystalline solid with a vapour pressure of about 10 mm. The complex can be readily sublimed. In Kei-F tubes small rhombus shaped plates grow slowly over a period of several days. At 30° crystal growth is rapid and well formed crystals can be obtained in 24 hr. These crystals tend to be twinned and produce overgrowths after several days. Sometimes faults and cracks develop after subjection to X-rays. Powder diffraction patterns show similarities to NO2WF7 and NO2TcFT[3], but the product is not iso-structural with these compounds. I.R. spectra of the vapour in equilibrium with solid NO2F.XeF+ show only bands of the component molecules indicating considerable or complete dissociation in the vapour. In the solid deposited on cold AgC1 windows at - 195°, absorptions at 585 and 2300 cm-1 are observed. Absorptions in nitryl fluoride complexes in the 2360-2400 c m - t region indicate the presence of the NO2 + cation; those around 570600 cm - 1 are indicative of Xe-F stretching modes or of the bending mode of the NO 2 + cation. The Raman spectrum of the solid shows an intense sharp band at 636 cm- 1, medium intensity bands at 562 and 524 cm- 1 and a number of weaker hands in the bending region. On the other hand, a weak Raman hand at 2305 cm- ~ may possibly be due to the antisymmetric stretching vibration of the nitryl cation resulting from breakdown of the selection rules in the solid. The symmetrical stretching vibration of NO2 + which should be Raman active at about 1400 cm- t appears to be extremely weak. This, and the fact that the i.r. allowed mode is lower than the normally occurring 2380 cm- 1 suggest incomplete fluoride ion transfer. Moreover, the symmetrical Xe-F stretching vibration at 636 cm- 1 is not shifted substantially from the corresponding one at 657 cm-1 in solid XeF e, whereas the symmetrical Xe-F stretching vibration in XeFs" is shifted down considerably to 540 cm-114, 12]. The latter ion has been definitely established in the solid state[12]. The complex NO2F.XeFe can probably not be formulated as the ionic form NO2 +XeFT-, but rather as an intermediate fluorine bridged structure. The complex NO2F.XeF6 hydrolyses rapidly to form a colourless solution. Although some gas bubbles are evolved during hydrolysis, the resulting solution possesses oxidizing properties characteristic of aqueous xenon(VI). When solutions+exposed to the atmosphere dry out, the residual solid is explosive. Although analogous complexes of XeFe are known in the salts CsXeF7 and RbXeFT[13], the NO2F adduct is the only one of this type which is volatile and amenable to vacuum manipulations.
Chemistry Department The University oJ Leicester Leicester LEI 7RH Department o[ Inorganic and Analytical Chemistry The Hebrew University Jerusalem Israel
JOHN H. HOLLOWAY
HENRY SELIG URI EL-GAD
12. S. W. Peterson, J. H. Holloway, B. A. Coyle and J. M. Williams, Science 173, 1238 (1971). 13. R. D. Peacock, H. Selig and I. Sheft, J. inorg, nucl. Chem. 28, 2561 (1966).