-3
$\begingroup$

The fluorine however it says my text book doesn't combine well with oxygen, nitrogen and itself but when about its oxides they are quite stable as $\ce{O3F2}$ has a melting point of $\pu{363 °C}$ means its quite stable.It also says that oxides of chlorine are unstable.

Now I have two questions:

  1. Why does Fluorine form stable oxide even though it's as electronegative as oxygen?
  2. Why does chlorine form unstable oxides?
$\endgroup$

closed as unclear what you're asking by aventurin, Tyberius, Mithoron, Gaurang Tandon, Nilay Ghosh Apr 3 '18 at 16:19

Please clarify your specific problem or add additional details to highlight exactly what you need. As it's currently written, it’s hard to tell exactly what you're asking. See the How to Ask page for help clarifying this question. If this question can be reworded to fit the rules in the help center, please edit the question.

  • 5
    $\begingroup$ Where did you even hear about something as unholy as O3F2 ?! Even if it does exist, it would be like most reactive thing on Earth! And by no means having such m.p. $\endgroup$ – Mithoron Apr 3 '18 at 14:46
  • 3
    $\begingroup$ Technically, fluorine does not form oxides. It forms oxygen fluorides. Fluorine tops oxygen in IUPAC's electronegativity pecking order. $\endgroup$ – Oscar Lanzi Apr 3 '18 at 14:46
  • 6
    $\begingroup$ Are you sure of that melting point? This book states that O3F2 melts at 84K, or -189 C. It also reports "mild explosions" in reactions of O3F2 at 90K, which is not the hallmark of a stable compound. $\endgroup$ – AndyW Apr 3 '18 at 14:56
4
$\begingroup$

Late edition:

To our surprise (most of us, evidently), the compound $\ce{O3F2}$ exist, at least in lower temperatures. The synthesis of $\ce{O3F2}$ was claimed by two Japanese scientists during World War II [Ref. 1], but their claims have not been accepted due to lack of quantitative analysis of the compound to prove the molecular formula. After about 20 years later, its existence was confirmed by a publication in the Journal of American Chemical Society in 1959 titled Ozone Fluoride or Trioxygen Difluoride, $\ce{O3F2}$ [Ref. 2], which gave detailed analysis of the compound. The abstract of the paper states that:

The existence of ozone fluoride, $\ce{O3F2}$, has been put on a solid foundation by isolating the pure compound and analyzing it. $\ce{O3F2}$ is a deep blood-red liquid, solidifying at $\pu{83 ^\circ K}$. and decomposing at about $\mathrm116^\circ$ or higher in a clean cut reaction to $\ce{O2}$ and $\ce{O2F2}$. It is an endothermic compound and is one of the most potent oxidizers known. It is more reactive than either $\ce{F2}$, $\ce{OF2}$ or mixtures of $\ce{O2}$ and $\ce{F2}$.

The description of physical properties states that:

It is a blood-red viscous liquid which can be refluxed and distilled with only slight decomposition in the range of $96$ to $\pu{114^\circ K}$ and at a pressure of $0.1$ to $\pu{1.5 mm}$. It remains liquid at $\pu{90^\circ K}$ and can thus be easily distinguished from $\ce{O2F2}$.

Note: $\ce{O2F2}$ is an orange solid discovered in 1933, which melts at $\pu{109.7^\circ K}$ to a red liquid. Thus, at $\pu{90^\circ K}$, it should be still a solid.

The paper also stated that, alike $\ce{O2F2}$, with heat evolution, $\ce{O3F2}$ also decomposes quantitatively to $\ce{O2F2}$ and $\ce{O2}$ at about $\pu{115^\circ K}$ ($\ce{2 O3F2 -> O2 + 2 O2F2}$).

I believe this reaction might be a one reason to ban Freon and other fluorinated compounds from common use to save the ozone layer. But its claimed melting point $\pu{363 °C}$ by Pakistan Zindabad is not quite right. According to the discoverer, the compound is not existing after $\pu{115^\circ K}$. At that temperature, it converts to $\ce{O2F2}$ (and $\ce{O2}$). At about $\pu{200^\circ K}$, $\ce{O2F2}$ will further dissociate quantitatively, to $\ce{O2 + F2}$, again under heat evolution.

References:

  1. Gmelin Handbook of Inorganic Chemistry: Fluorine - Compounds with Oxygen and Nitrogen; Suppliment Vol 4, 8th Edn., Susanne Jager, et al., Eds., Springer-Verlag: Berlin, Germany, 1986, p. 103-104 (Chapter 3. Compounds of Fluorine: Fluorine and Oxygen: 3.13. Trioxygen Difluoride, $\ce{O3F2}$).

  2. Ozone Fluoride or Trioxygen Difluoride, $\ce{O3F2}^1$: A. D. Kirshenbaum, A. V. Grosse, J. Am. Chem. Soc., 1959, 81(6), 1277–1279 (https://pubs.acs.org/doi/abs/10.1021/ja01515a003).

$\endgroup$

Not the answer you're looking for? Browse other questions tagged or ask your own question.