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The noble gases, belonging to the ZERO GROUP of the Periodic Table were considered totally inert, chemically; this is still true only for HELIUM (He), NEON (Ne) and ARGON (Ar), but not anymore for the heaviest of the group, KRIPTON (Kr), XENON (Xe) and the radioactive RADON (Rn).
So, we can say today that the noble gases are not chemically inert, but only that they react with difficulty.

Their reactivity was shown in 1962 by N. Bartlett who had reached to produce a salt containing the ion O2+ , the Dioxigenyl Hexafluoroplatinate V, whose formula O2+ PtF6-.
Given that the ionization potential of oxygen (O2 -> O2+ + e-) was 12.2 eV, very similar to that of Xe (12.13 eV), he thought that the same reaction could occur using Xe at the place of O2:

Xe + PtF6 -> Xe+ PtF6-

He was right and he produced the HEXAFLUOROPLATINATE (V) OF XENON (XePtF6) a stable orange coloured solid that was the first compound known of a noble gas.
In the following years, other compounds were prepared of Kr (only the fluoride KrF2, decomposing at 0C) and much more of Xe with F, O, more stable, like XeF2, XeF4, XeF6, XeO3, Xe(OH)4 and others with Cl and N, much less stable.
Since 1968, were prepared the oxigenated salts of Xe with alkaline metals, like RbFXeO3 (Rubidium fluoroxenate VI), that are stable until 200-300 C.
The energies of the chemical bonds of Xe and Kr are:
- Xe-F: 125 kj/mole
- Xe-O: 85 "
- Kr-F: 50 "
So, we can explain why the compounds of Kr are less stable also with fluorine.

CHEMICAL REACTIONS
I report some informations about the chemistry of Xe with F and O, the most important, being the only to give stable compounds.

- XeF4: It's formed making react Xe with an excess of F2 at about 400 C and 5 Atm. If we quickly cool at -50 C, we can obtain a mix of XeF2, in higher percentual and XeF4, in minor amount.
If, instead, we make the same reaction at 50 Atm, we obtain XeF6,
All these fluorides are solid and their melting temperatures are
129 C for XeF2, 115 C for XeF4 and 50 C for XeF6.
All the Xe fluorides are strongly oxidant:

XeF2 + 2 HI -> 2 HF + I2 + Xe
XeF4 + 4Hg -> 2 Hg2F2 + Xe

and in water they oxidate it:

2 XeF2 + 2 H2O -> 4 HF + O2 + 2 Xe
XeF4 + 2 H2O -> 4 HF + O2 + Xe
2 XeF6 + H2O -> 12 HF + 3 O2 + 2 Xe

In reality, these fluorides gives also other parallel reaction with water with production of XeO3 (xenon trioxide):

3 XeF4 + 6 H2O -> 2 XeO3 + 12 HF + Xe

from XeO3, all other Xe compounds of oxygen are derived.

- XeO3:
In alkaline solution, it forms the XENATES (HXeO-) that give a further DISPROPORTION reaction producing Xe and PERXENATES, with oxidation number (ON) of +8 for Xe and very strong oxidants
:

XeO3 + NaOH -> NaHXeO4
4 NaHXeO4 + 8 NaOH -> 3 Na4XeO6 + 6 H2O + Xe

Using NaOH the salt formed is insoluble and precipitates as
Na4XeO6*8H2O that treated with H2SO4 (sulfuric acid) gives XeO4, gaseous and unstable:

Na4XeO6 + 2 H2SO4 -> 2 NaSO4 + 2 H2O + XeO4

So, all these compounds of Xe with ON = +2, +4, +6 and +8 are all originated from the fluorides of this element.
From perxenates many OXYFLUORIDES can be prepared:

XeOF4, XeO2F2, XeO2F4, XeO3F2, and so on

Learn more about this author, Aldo Bonincontro.
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